IRLF 


SB    35    373 


Edmund  O'Neill 


ELEMENTS 


OF 


GENERAL  CHEMISTRY 


WITH 


EXPERIMENTS. 


BY 

JOHN   H.  LONG,  M.  S.,  Sc.  D., 

PROFESSOR   OF   CHEMISTRY   AND    DIRECTOR    OF    THE    CHEMICAL    LABORATORIES   IN    THE   SCHOOLS   OF 
MEDICINE   AND   PHARMACY   OF    NORTHWESTERN    UNIVERSITY. 


CHICAGO  : 
E.   H.  COLEGROVE, 

SALES    AGENT. 

1898. 


Entered  according  to  act  of  Congress,  in  the  year  1898, 

BY  JOHN  H.  LONG, 
in  the  office  of  the  Librarian  of  Congress,  at  Washington. 


INMEMORIAM, 


PREFACE. 


In  the  preparation  of  this  book  the  author  has  used  a 
part  of  his  work,  "  Experimental  and  Analytical  Chemis- 
try," now  out  of  print,  and  has  added  enough  new  matter 
to  make  of  it  a  complete  text-book  of  elementary  general 
chemistry,  sufficient  for  the  wants  of  college  students 
beginning  the  subject. 

In  too  many  instances  the  student  is  introduced  to 
qualitative  analysis  as  his  first  laboratory  work,  and  this  is 
followed  by  gravimetric  analysis  to  complete  a  course. 
This  plan  certainly  gives  the  beginner  a  distorted  idea  of 
the  relative  importance  of  analytical  chemistry  in  the  study 
of  the  science;  for  the  beginner  a  knowledge  of  the  proper- 
ties of  a  substance,  the  methods  of  its  preparation  and  its 
uses  is  far  more  important  than  acquaintance  with  methods 
of  separation,  and  general  illustrative  experiments  should, 
therefore,  be  made  the  foundation  work  in  the  laboratory. 

It  is  the  belief  of  the  author  that  much  that  is  demon- 
strated by  the  teacher  in  the  classroom  may  profitably  be 
repeated  by  the  student  in  the  laboratory.  Repetition  is 
necessary  to  fix  elementary  principles  thoroughly  in  the 
mind  of  the  beginner.  The  list  of  experiments  here  offered 
embraces  the  work  of  this  character  which  has  been 
required  in  the  author's  classes  during  the  past  ten  years. 
Most  of  these  exercises  are  simple  and  easily  performed; 
others  are  longer  or  more  complex,  and  are  therefore 
described  in  considerable  detail,  but  all  of  them  may  be 
performed  by  the  aid  of  comparatively  simple  apparatus, 


889793 


and  all,  it    is  believed,  illustrate  important  facts  or  prin- 
ciples. 

In  the  descriptive  part  of  the  book  the  author  has  kept 
in  mind  the  fact  that  it  is  intended  for  beginners,  few  of 
whom  expect  to  become  specialists  in  chemistry,  and  he  has, 
therefore,  made  the  presentation  of  matter  as  practical  as 
possible.  Some  important  substances  and  technical 
processes  are  described  more  fully  than  is  usually  thought 
necessary  in  an  elementary  book.  No  chemical  theory  is 
introduced  in  the  earlier  chapters,  but  after  the  student  has 
been  made  familiar  with  important  principles  by  experi- 
ment it  is  gradually  presented.  In  explaining  the  atomic 
theory  an  attempt  has  been  made  to  show  in  a  very  ele- 
mentary manner  the  important  steps,  historically,  in  its 
development.  It  is  believed  that  this  method  will  give  the 
student  the  clearest  insight  into  a  subject  which  is,  at  best, 
hard  to  grasp  and  which  is  seldom  mastered. 

The  author  wishes  to  acknowledge  the  very  valuable 
assistance  rendered  him  by  his  friend  and  colleague, 
Dr.  Charles  H.  Miller,  in  reading  proofs  and  in  other  ways 
helping  in  the  publication  of  the  book. 

THE  AUTHOR. 

CHICAGO,  1898. 


TABLE  OF  CONTENTS. 


CHAPTER  I.     Introductory I    1 

CHAPTER  II.      Oxygen,  Hydrogen  and  their  Compounds  30 

CHAPTER  III.  Chlorine  and  Hydrochloric  Acid. — 
Theoretical  Considerations 66 

CHAPTER  IV.  Compounds  of  Chlorine  with  Oxygen. — 
Bromine,  Iodine,  Fluorine  and  their  Compounds..  90 

CHAPTER  V.  Nitrogen  and  the  Atmosphere. —  Gas 
Problems 106 

CHAPTER  VI.     Compounds  of  Nitrogen 120 

CHAPTER  VII.  Sulphur  and  its  Compounds,  Selenium 
and  Tellurium 146 

CHAPTER  VIII.  Silicon  and  Boron  and  their  Com- 
pounds  172 

CHAPTER  IX.  Phosphorus  and  Arsenic  and  their  Com- 
pounds  183 

CHAPTER  X.  Carbon  and  some  of  its  important  Com- 
pounds   202 

CHAPTER  XL     Atomic  and  Molecular  Weights 231 

CHAPTER  XII.  Classification  of  the  Elements.  Gen- 
eral Properties  of  the  Metals  and  their  Salts 249 

CHAPTER  XIII.  The  Alkali  Metals:  Lithium,  Sodium, 
Potassium,  Rubidium  and  Caesium. — Ammonium 
Compounds 271 


CHAPTER  XIV.  The  Copper  Group:  Copper,  Silver 
and  Gold 288 

CHAPTER  XV.  The  Alkali  Earth  Group:  Beryllium, 
Magnesium,  Calcium,  Strontium  and  Barium. — The 
Spectroscope 310 

CHAPTER  XVI.     Zinc,  Cadmium  and  Mercury 328 

CHAPTER  XVII.  Boron,  Aluminum,  Gallium,  Indium, 
Thallium,  Scandium,  Yttrium,  Lanthanum  and 
Ytterbium 330 

CHAPTER  XVIII.  The  Carbon  Group:  Carbon,  Silicon, 
Germanium,  Tin  and  Lead. — The  Titanium  Group: 
Titanium,  Zirconium,  Cerium  and  Thorium 344 

CHAPTER  XIX.  The  Nitrogen  Group:  Nitrogen,  Phos- 
phorus, Vanadium,  Arsenic,  Columbium,  Antimony, 
Tantalum  and  Bismuth 354 

CHAPTER  XX.  The  Chromium  Group:  Chromium, 
Molybdenum,  Tungsten  and  Uranium.  Relations 
to  the  Oxygen  Group 362 

CHAPTER  XXI.  Manganese  and  its  Relations  to  the 

Halogen  Group 369 

CHAPTER  XXII.  The  Iron  Group:  Iron,  Nickel  and 

Cobalt 375 

CHAPTER  XXIII.  The  Platinum  Group:  Ruthenium, 
Rhodium,  Palladium,  Osmium,  Indium  and  Plati- 
num.. ..389 


CHAPTER  I. 


INTRODUCTORY. 

IN  BEGINNING  the  study  of  chemistry  in  the  labora- 
tory or  classroom  the  student  should  learn  to  consider 
each  experiment  performed  as  a;  question ,  and  th&  result 
obtained  its  answer.  Chemistry  is^fe^mitieritlyran^exper- 
imental  science  in  which  matte*  under  A^t^in,  conditions 
is  the  subject  of  investigation. 

By  experiment  and  observation  we  seek  to  determine 
the  properties  of  this  matter,  to  divide  it  into  groups,  to 
analyze  and  decide  what  is  simple  and  what  compound,  to 
find  the  action  which  one  kind  of  matter  exerts  upon 
another  and  how  each  one  behaves  under  the  influence  of 
heat,  light,  electricity  and  other  forces.  We  seek  also  to 
find  the  simplest  and  best  means  of  producing  different 
kinds  of  matter,  and  to  discover  tests  by  which  they  may 
be  always  recognized. 

This  knowledge,  with  more  to  be  acquired,  when  prop- 
erly classified  and  arranged  in  a  consistent  system,  consti- 
tutes the  science  of  chemistry.  The  beginner  can  best 
obtain  acquaintance  with  this  science  by  his  own  experi- 
ments in  the  laboratory  under  the  guidance  of  an  instruct- 
or. Much  can  be  and  must  be  learned  from  books,  it  is 
true,  but  the  knowledge  which  is  most  satisfactory  and 
most  lasting  when  acquired  is  that  which  the  student 
gathers  by  direct  contact  with  the  thing  under  study. 

In  many  lines  observation  alone  brings  but  limited 
knowledge.  For  instance,  of  the  air  or  of  the  water  every- 
where around  us,  we  would  know  indeed  but  little  if  un- 
aided by  experiment.  When  we 'make  an  experiment  on 
an  object  we  take  the  thing,  in  a  sense,  within  our  grasp 
and  look  at  it  from  different  sides,  placing  it  under  new 
and  varied  conditions,  and  by  so  doing  learn  many  of  its 


2  GENERAL  CHEMISTRY. 

important  qualities  and  peculiarities.  Asking  ourselves 
how  it  would  behave  under  certain  conditions,  we  make 
the  experiment  and  find  out.  In  chemistry  we  study  mat- 
ter as  undergoing  change. 

We  are  acquainted  with  matter  in  three  general  forms 
or  conditions,  the  gaseous,  liquid  and  solid  states,  and  we 
shall  first  give  our  attention  to  a  brief  consideration  of 
these. 

The  Three  States  of  Matter. 

Many  kinds  of  matter  are  found  to  exist  in  the  three 
forms  mentioned,,  but  k>r  each  substance  there  is  a  condi- 
tion in:  which'  it  is^mos-t  stable  and  most  usually  found. 
Th£  charge:  from;  one  condition  to  the  others  is  generally 
most  r^raiiiit^'broti^htalDOutby  a  change  of  temperature;  a 
low  temperature  being  favorable  to  the  maintenance  of  the 
solid  condition  while  a  high  temperature  aids  in  the  forma- 
tion of  gases  or  vapors.  We  have  in  water  a  familiar 
illustration  of  a  substance  well  known  in  the  three  condi- 
tions, but  many  other  common  substances  can  readily  be 
made  to  pass  from  one  of  these  conditions  to  the  others,  as 
can  be  shown  by  experiment. 

Ex.  I.  Let  the  student  apply  heat  to  a  test-tube  one-third  filled 
with  sulphur.  At  a  temperature  of  about  115°  C.  it  melts  to  a  yellow 
liquid  which  grows  darker  by  application  of  more  heat  and  becomes 
viscid.  At  a  still  higher  temperature  the  viscid  mass  becomes  thinner, 
and  finally  boils  at  a  temperature  of  about  450°  C.  Application  is  made 
of  this  fact  in  the  refining  of  sulphur  by  distillation. 

In  this  experiment  the  vapor  of  the  sulphur  usually  ignites  at  the 
mouth  of  the  test-tube  and  burns  with  a  pale  blue  flame,  forming  sul- 
phurous oxide,  as  will  be  explained  later. 

Ex.  2.  In  a  somewhat  narrow  test-tube  melt  two  or  three  grams 
of  camphor.  This  passes  from  the  solid  to  the  liquid  condition  at  a 
temperature  of  175°  and  boils  at  204°.  Vapors  are  given  off  even  at  low 
temperatures,  from  which  it  follows  that  in  experimenting  with  quite 
small  pieces  of  camphor  the  middle  or  liquid  condition  may  be  over- 
looked If  the  tube  taken  is  long  enough,  10  to  12  Cm.,  the  vapor  from 
the  boiling  liquid  will  condense  on  the  upper  and  cooler  part. 

With  iodine  and  several  other  bodies  the  phenomena  of 
vaporization  are  very  similar.  Iodine  melts  at  about  115°, 
but  gives  off  vapors  at  a  lower  temperature.  So  rapid  is 


GENERAL  CHEMISTRY.  3 

vaporization  above  the  melting  point,  that  the  temperature 
of  actual  ebullition  cannot  be  accurately  observed.  It  is 
above  200°. 

With  ammonium  chloride,  or  sal  ammoniac,  the  behav- 
ior is  different.  The  substance  gives  off  no  vapors  at  the 
ordinary  temperature,  but  readily  at  high  temperatures. 

Ex.  3.  In  a  test-tube  heat  some  of  the  ammonium  chloride  in  fine 
crystals.  Observe  that  it  does  not  melt,  but  at  a  sufficiently  high  tem- 
perature gives  off  dense  white  vapors,  which  soon  condense  to  crystal- 
line grains.  We  have  here  the  passage  from  the  first  to  the  third  state 
without  liquefaction. 

Many  of  our  common  and  best  known  substances  can- 
not be  obtained  in  the  state  of  vapor,  and  some  not  even 
in  the  liquid  condition,  because  they  surfer  decomposition 
when  strongly  heated. 

The  red  oxide  of  mercury  is  a  good  illustration  of  this, 
as  it  separates  into  mercury  and  oxygen  by  heat. 

Common  limestone  breaks  up  when  heated,  yielding 
quicklime  and  carbonic  acid  gas.  Common  salt  and 
potassium  chlorate  may  be  liquefied,  but  suffer  decomposi- 
tion when  heated  to  higher  temperatures. 

The  so-called  organic  substances  are  those  in  which 
passage  through  the  three  states  can  be  most  readily  ob- 
served, as  the  temperature  of  vaporization  is  in  general 
much  lower  here  than  among  the  inorganic  compounds. 
The  change  of  all  bodies  from  the  liquid  condition  to  that 
of  a  gas  or  vapor  depends  not  only  on  temperature,  but 
also  on  the  pressure  on  it,  that  of  the  atmosphere  usually. 
Variations  of  the  atmospheric  pressure  cause  a  change  in 
the  temperature  to  which  a  substance  must  be  brought  to 
change  it  from  a  liquid  to  a  vapor. 

Vaporization  follows  at  a  lower  temperature  by  de- 
crease of  the  air  pressure  on  the  heated  liquid.  Hence  it 
is  that  many  substances  which  cannot  be  distilled  under 
the  ordinary  atmospheric  pressure  without  decomposition 
can  be  easily  and  safely  distilled  in  a  partial  vacuum. 

Solutions. 

When  common  salt  is  thrown  in  water  and  the  mixture 
stirred  the  salt  gradually  disappears,  leaving  finally  a  clear 


4  GENERAL  CHEMISTRY. 

liquid  which  in  appearance  cannot  be  distinguished  from 
the  water.  Many  other  substances  behave  in  the  same 
manner,  for  instance,  sugar,  saltpeter,  soda,  borax  and  sal 
ammoniac.  We  apply  the  name  solution  to  the  clear  mix- 
tures of  these  substances  with  water.  The  sugar,  salt  and 
other  substances  are  said  to  be  soluble  in  water.  They  are 
soluble,  also,  in  other  liquids.  In  such  cases  the  particles 
of  the  solid  seem  to  distribute  themselves  among  the  liquid 
particles,  and  in  every  instance  there  is  a  limit  to  the 
power  or  capacity  of  the  water  for  dissolving  the  solid. 

In  the  illustrations  given,  as  in  manyothers,  the  solvent 
and  body  dissolved  exert  no  decomposing  action  on  each 
other,  because  the  two  can  be  readily  separated  and 
obtained  in  their  original  conditions.  To  make  this  point 
plain  let  the  student  make  the  following  experiment : 

Ex.  4.  Into  some  distilled  water  stir  clean,  pure  salt,  a  little  at  a 
time,  until  the  water  becomes  saturated^  that  is,  until  it  will  no  longer 
take  up  any  more  of  the  salt.  In  50  cubic  centimeters  at  the  ordinary 
temperature  we  can  dissolve  in  this  manner  about  18  grams  of  salt. 

Next  pour  about  half  of  the  solution  into  an  evaporating  dish,  place 
this  on  wire  gauze  or  on  a  sand-bath  (sand  in  an  iron  dish)  over  the  low 
gas  flame  from  a  Bunsen  burner  and  heat  slowly.  The  water  gradually 
disappears  or  passes  off  in  the  form  of  vapor,  leaving  at  last  the  dis- 
solved salt  as  a  clear  white  crystalline  mass. 

Many  substances  not  soluble  in  water  are  soluble  in 
alcohol,  ether,  chloroform  or  other  liquid,  and  usually 
without  change;  that  is,  by  evaporation  of  the  menstruum 
the  substance  may  be  recovered  as  was  the  salt  in  the 
above  experiment.  But  other  bodies  dissolve  only  by 
decomposition.  For  example,  marble  is  not  soluble  in 
water,  but  it  can  be  quickly  dissolved  by  action  of  certain 
acids. 

Ex.  5.  In  a  small  beaker  take  a  few  grams  of  chalk  or  powdered 
marble  (commonly  called  marble  dust).  Add  water  and  stir  or  shake 
the  mixture  thoroughly.  Then  allow  to  settle,  and,  as  far  as  can  be  de- 
termined by  the  eye,  it  will  be  noticed  that  the  marble  remains  undis- 
solved.  Next  add,  a  few  drops  at  a  time,  some  hydrochloric  acid  and 
the  escape  of  gas  which  follows  shows  that  an  important  change  is 
taking  place.  Gradually  add  more  acid  until  the  effervescence,  after 
shaking,  ceases.  There  should  be  left  now  a  clear  or  nearly  clear  liquid 
or  solution,  and  by  evaporating  this  in  a  small  dish  or  beaker,  the  slight 
excess  of  acid  employed  in  making  it  will  be  driven  rff.  What  remains 
is  soluble  in  water,  while  the  original  marble  was  not. 


GENERAL  CHEMISTRY.  5 

The  action  of  the  acid  here  has  been  to  convert  an  in- 
soluble body  into  one  readily  soluble  in  water.  We  can- 
not properly  speak  of  the  solution  as  a  solution  of  marble, 
as  this  substance  is  no  longer  present.  As  another  simple 
illustration  of  solution  effected  by  conversion  into  a  new 
substance,  the  action  of  acids  on  many  metals  may  be 
referred  to.  It  will  be  shown  later  that  iron,  zinc  and 
other  common  metals  dissolve  readily  in  hydrochloric  or 
sulphuric  acid.  During  the  action  of  the  acid  on  the  metal 
a  gas  escapes  and  there  is  left  dissolved  a  combination, 
termed  a  salt,  of  a  part  of  the  acid  with  the  metal. 

In  a  mixture  holding  a  body  in  solution  and  something 
insoluble  in  suspension,  the  latter  may  be  separated  by 
filtration,  that  is,  by  passing  the  liquid  through  an  appara- 
tus termed  a  filter,  which  holds  the  particles  not  actually 
in  solution.  As  a  filtering  medium,  paper,  sand,  porous 
stone,  felt,  and  other  substances  may  be  used.  In  illus- 
tration of  this,  make  the  following  experiment: 

Ex.  6.  In  a  beaker  mix  some  common  salt  and  clean  marble  dust. 
Pour  on  water  and  shake  thoroughly.  Allow  to  subside  and  then  pour 
the  liquid  on  a  paper  filter.  (For  method  of  making  a  paper  filter  the 
instructor  must  be  consulted.)  To  the  residue  in  the  beaker  add  more 
water,  stir  again  and  pour  through  the  same  filter.  Finally,  wash  the 
residue  itself  from  the  beaker  onto  the  filter  and  allow  it  to  drain. 
When  dry  it  will  be  recognized  as  the  original  marble  dust.  The  liquid 
which  passed  through  the  filter,  or  filtrate,  on  evaporation  yields  the 
salt. 

In  this  experiment  the  fine  pores  or  openings  in  the 
paper  permit  the  passage  of  the  liquid  and  the  salt  dis- 
solved in  it,  but  not  the  passage  of  even  the  finest  parti- 
cles of  the  undissolved  marble  dust.  Filtration  is  one  of 
the  most  common  operations  of  analytical  chemistry. 

Bodies  differ,  when  soluble,  very  greatly  in  the  extent 
of  their  solubility.  While  common  salt  will  dissolve  in 
less  than  three  times  its  weight  of  water,  at  the  ordmary 
temperature,  or  cane  sugar  in  about  half  its  weight  of 
water,  there  are  required  for  gypsum  nearly  400  and  for 
morphine  nearly  1,000  parts  of  water.  Indeed,  it  may  be 
said  that  no  substance  is  absolutely  insoluble  in  water, 
but  where  the  degree  of  solubility  is  so  small  that  several 


6  GENERAL  CHEMISTRY. 

thousand  parts  of  water  are  required  for  solution  it  is  cus- 
tomary to  speak  of  the  body  in  question  as  practically 
insoluble.  In  many  classes  of  investigations,  however, 
even  very  slight  degrees  of  solubility  must  be  taken  into 
consideration. 

The  temperature  of  a  liquid  has  in  most  cases  a  marked 
influence  on  its  solvent  power.  It  usually  happens  that 
the  solubility  of  a  substance  is  increased  by  increase  of 
temperature,  but  this  is  not  always  the  case.  We  know  of 
a  few  common  substances  which  are  actually  less  soluble 
in  warm  liquids  than  they  are  in  cold.  A  striking  illustra- 
tion of  the  change  of  solubility  with  change  of  tempera- 
ture is  shown  by  the  behavior  of  niter  or  saltpeter  with 
water.  At  a  temperature  of  20°  C,  that  is  at  a  common 
room  temperature,  100  parts  by  weight  of  water  dissolve 
about  31  parts  of  the  niter,  but  at  a  temperature  of  100°  C., 
that  is  at  the  temperature  of  boiling  water,  nearly  250 
parts  may  be  dissolved.  The  following  experiment  may 
be  made  by  way  of  illustration: 

Ex.  7.  In  a  test-tube  take  about  10  Cc.  of  water  at  the  laboratory 
temperature.  Add  to  it  some  powdered  saltpeter,  a  little  at  a  time, 
close  with  the  thumb  and  shake  after  each  addition  until  the  water  be- 
comes saturated,  that  is  until  it  will  no  longer  dissolve  the  added  salt- 
peter. It  will  be  observed  that  as  the  solid  goes  into  solution  the 
temperature  of  the  liquid  becomes  lower.  Now  gradually  heat  the 
solution  in  the  Bunsen  burner  flame  and  from  time  to  time  add  more  of 
the  saltpeter.  In  this  way  it  will  be  seen  that  a  clear  solution  may  be 
made  containing  many  times  the  weight  of  the  substance  dissolved  in 
the  cold.  When  it  has  become  saturated  at  the  boiling  heat  set  it  aside 
to  cool  slowly.  After  a  time  the  test-tube  will  be  found  to  contain  a 
mass  of  crystals  deposited  by  the  cooling  of  the  liquid. 

Other  interesting  examples  of  the  same  effect  of  tem- 
perature may  be  referred  to.  At  0°,  100  Gm.  of  water  dis- 
solves about  4  Gm.  of  crystallized  potassium  alum,  but  at 
100°  the  same  weight  of  water  dissolves  over  350  Gm.  of 
the  alum.  At  0°,  100  Gm.  of  water  dissolves  about  5  Gm. 
of  crystallized  oxalic  acid,  but  at  100°  nearly  350  Gm.  may 
be  dissolved. 

A  solution  is  said  to  be  saturated  at  a  certain  temper- 
ature when  it  contains  as  much  of  a  substance  as  it  will  hold 
at  this  temperature.  It  is  a  peculiarity  of  many  substances, 


GENERAL  CHEMISTRY.  7 

however,  that  they  may  be  temporarily  dissolved  in  water 
or  other  liquid  in  amount  greater  than  can  be  permanently 
held  at  the  same  temperature.  A  solution  so  produced  is 
said  to  be  supersaturated.  This  condition  is  most  readily 
attained  by  dissolving  a  salt  by  the  aid  of  heat  until  the 
solution  becomes  saturated  at  a  high  temperature.  On 
carefully  pouring  off  some  of  the  clear  hot  liquid  it  may 
often  be  cooled  to  the  air  temperature  and  kept  a  long  time 
without  precipitation.  Let,  however,  a  small  crystal  of 
the  dissolved  substance  be  dropped  into  the  cool  liquid,  a 
precipitate  from  the  solution  will  appear  and  settle  out 
until  the  amount  remaining  dissolved  is  just  sufficient  to 
constitute  a  normally  saturated  solution  at  the  given  low 
temperature.  The  change  from  the  state  of  supersatura- 
tion  to  that  of  normal  saturation  is  here  brought  about  by 
addition  of  some  of  the  same  salt  that  is  dissolved,  and  in 
general,  supersaturation  in  a  solution  at  a  given  tempera- 
ture may  be  detected  in  this  manner.  It  may  be  detected 
also  by  other  means.  The  phenomenon  is  one  of  so  much 
importance  that  it  will  be  illustrated  by  experiment. 

Ex.  8.  In  each  of  three  perfectly  clean  beakers  holding  about  250 
Cc.  dissolve  50  Gm.  of  pure,  powdered,  crystallized  sodium  sulphate  in 
25  Cc.  of  distilled  water.  Heat  to  30°-35°  to  complete  the  solution. 
When  a  clear  solution  is  obtained  cover  each  one  of  the  beakers  with  a 
piece  of  paper  and  set  them  aside  to  cool  in  a  perfectly  still  place. 
When  the  three  solutions  are  quite  cool  test  them  as  follows:  Remove 
the  papers  and  into  one  beaker  drop  a  small  crystal  of  the  sodium  sul- 
phate. By  means  of  a  glass  rod  rub  the  bottom  and  sides  of  the  second 
beaker,  while  the  contents  of  the  third  maybe  poured  out  into  a  dry 
beaker.  In  each  case  the  equilibrium  of  the  solution  is  destroyed  and  a 
precipitate  of  the  dissolved  salt  settles  out. 

The  substances  which  most  readily  form  supersaturated 
solutions  are  those  which  combine  with  large  amounts  of 
"water  of  crystallization."  This  term  will  be  explained 
later.  The  alums  and  borax  resemble  sodium  sulphate  in 
this  behavior. 

Crystallization. 

In  the  foregoing  the  formation  of  solutions  has  been 
explained  and  it  has  been  shown  that  the  dissolved  sub- 
stance may  often  be  readily  separated  or  recovered  from 


8  GENERAL  CHEMISTRY. 

the  solution.  In  the  case  of  bodies  not  decomposed  by  the 
menstruum  this  is  most  readily  effected  by  evaporation,  as 
illustrated  by  the  recovery  of  salt  or  sugar  dissolved  in 
water.  The  widest  application  of  this  fact  is  made  in  the 
arts  in  the  production  or  purification  of  numerous  impor- 
tant substances.  Concentration  of  a  solution  is  often  suf- 
ficient to  throw  out  the  dissolved  substance. 

The  solid  very  frequently  assumes  what  is  termed  the 
crystalline  form  as  it  leaves  the  solution,  and  this  is  gener- 
ally the  case  when  it  is  deposited  slowly,  as  by  the  gradual 
cooling  of  a  liquid.  In  the  experiment  with  saltpeter  the 
formation  of  crystals  was  shown,  but  the  phenomenon  can 
be  better  illustrated  by  the  use  of  another  substance. 

Ex.  9.  Dissolve  25  Gm.  of  powdered  alum  in  75  Cc.  of  water  by 
aid  of  heat.  Filter  the  hot  solution  into  a  clean  beaker  which  may  then 
be  set  aside  in  a  quiet  place  for  spontaneous  evaporation.  Several 
hours  or  over  night  should  be  given  for  this  and  at  the  end  of  the  time 
large  crystals  of  alum  will  be  found.  To  prevent  the  too  rapid  cooling 
of  the  solution  immediately  after  its  preparation  the  vessel  containing  it 
may  be  wrapped  in  cotton,  or  better  still,  in  felt.  Slow  cooling  favors 
the  production  of  large,  well-formed  crystals. 

Ex.  10.  Prepare  strong  solutions  of  copper  sulphate,  or  blue 
vitriol,  and  chrome  alum  by  dissolving  about  15  Gm.  of  each  in  a  small 
quantity  of  warm  water.  For  each  about  50  Cc.  of  water  should  be 
used.  Filter  the  solutions  into  small  clean  beakers,  which  set  aside  in 
a  quiet  place  protected  from  dust  for  several  days.  The  dissolved  salts 
will  begin  after  a  time  to  crystallize  out  as  the  solvent  water  evaporates 
spontaneously.  Slow  evaporation  and  a  low  temperature  favor  here,  as 
before,  the  formation  of  large  crystals. 

Very  perfect  crystals  of  many  substances  may  be  made 
by  a  slight  modification  of  the  above  experiment.  If  in 
very  strong  solutions  of  the  alums,  blue  vitriol,  potassium 
dichromate  or  potassium  ferrocyanide,  for  instance,  a 
small  crystal  of  the  same  substance  be  suspended  by  means 
of  a  fine  thread  this  crystal  will  serve  as  a  nucleus  around 
which  a  deposit  forms  as  the  solutions  become  concen- 
trated by  spontaneous  evaporation.  A  very  pretty  effect 
is  obtained  by  growing  in  this  manner  a  good  crystal  of 
chrome  alum.  This  is  then  suspended  in  a  cold  satu- 
rated and  clear  solution  of  common  potash  alum,  when  the 
growth  continues,  the  potash  alum  being  deposited  on  the 


GENERAL  CHEMISTRY.  9 

chrome  alum.  Many  substances  which  are  isomorphous, 
that  is,  have  the  same  crystalline  form,  can  be  crystallized 
together  in  this  manner. 

The  process  of  crystallization  is  employed  in  many  in- 
dustries and  in  chemical  investigation  on  the  small  scale 
for  the  purification  of  substances.  This  can  be  illustrated 
by  an  experiment. 

Ex.  ii.  Dissolve  some  crude  common  salt  in  hot  water  to  make  a 
saturated  solution.  Filter  this  hot  into  an  evaporating  dish  and  con- 
centrate a  little  by  heat.  On  allowing  now  to  cool,  some  of  the  salt  will 
separate  in  pure  white  form.  By  repeating  the  operation  on  the  liquid 
remaining,  the  mother  liquor,  it  is  called,  further  crops  of  crystals  may 
be  obtained. 

As  salt  is  found  in  nature  its  natural  contaminations 
are  usually  substances  much  more  soluble  than  it  is. 
These  are  therefore  left  in  the  mother  liquor.  The  first 
crops  of  crystals  are  the  purest,  but  if  the  concentration 
be  carried  too  far,  the  salt  obtained  may  be  mixed  with 
these  impurities.  The  filtration  at  the  beginning  of  the 
operation  above  was  intended  to  remove  insoluble  sub- 
stances only. 

By  fractional  crystallization  it  is  often  possible  to  sepa- 
rate two  or  more  substances  from  a  solution.  This  is  true 
where  the  substances  dissolved  differ  greatly  in  their  de- 
grees of  solubility.  Sodium  chloride  may  be  separated 
from  sodium  nitrate  in  this  manner,  and  copper  sulphate 
from  potassium  sulphate.  In  concentrating  solutions  of 
either  one  of  these  mixtures  the  least  soluble  substance 
will  begin  to  separate  first.  The  first  fraction  may  be  very 
nearly  pure.  By  continuing  the  concentration  and  crystal- 
lization from  the  mother  liquors  the  last  fractions  obtained 
may  be  nearly  pure  crystals  of  the  most  soluble  substance. 
By  dissolving  now,  the  first  fraction  obtained  in  water  and 
crystallizing  again  the  crop  of  crystals  obtained  will  be 
nearly  or  quite  pure  in  some  cases.  The  mother  liquor  is 
used  as  the  solvent  for  the  second  fraction  which  yields 
now  a  fresh  portion  of  the  least  soluble  salt  and  holds  more 
of  the  most  soluble.  By  a  continuation  of  this  method  the 
most  soluble  constituent  can  be  concentrated  in  a  solution 
practically  free  from  the  others,  and  then  crystallized  itself. 


10  GENERAL  CHEMISTRY. 

Water  of  Crystallization. 

Many  substances  in  crystallizing  from  aqueous  solution 
unite  with  a  part  of  the  water,  holding  it  in  the  form  de- 
scribed as  water  of  crystallization.  Other  substances  crys- 
tallize in  the  anhydrous  form — that  is,  they  hold  no  water. 
Common  salt  and  saltpeter  are  familiar  illustrations  of 
bodies  belonging  to  the  second  group,  while  blue  vitriol, 
alum  and  Glauber's  salt  are  common  substances  which 
contain  water  of  crystallization.  Blue  vitriol  is  a  combina- 
tion of  copper  sulphate  with  water,  and  when  the  sub- 
stance is  powdered  and  strongly  heated  the  water  is  driven 
off,  leaving  the  copper  sulphate  in  the  anhydrous  or  pure 
form.  This  pure  copper  sulphate  is  no  longer  blue,  but 
white.  In  ordinary  usage  the  term  copper  sulphate  is 
understood  to  refer  to  the  common  blue  crystallized  com- 
pound. The  behavior  of  this  substance  when  heated  may 
be  shown  by  experiment. 

Ex.  12.  In  a  narrow  test-tube  heat  a  few  grams  of  powdered  blue 
vitriol  in  the  gas  flame,  but  not  to  a  high  temperature.  To  avoid  too 
great  heating  the  tube  may  be  moved  backward  and  forward,  and 
turned  meanwhile  between  the  fingers,  at  a  point  some  distance  above 
the  hottest  part  of  the  flame.  Four-fifths  of  the  water  held  by  the  sub- 
stance is  given  off  at  a  temperature  not  far  from  100°  C.f  and  may  be 
oeen  as  vapor  in  the  tube.  The  sulphate  is  left  as  a  bluish  white  pow- 
der after  the  escape  of  the  vapor.  The  heat  may  now  be  increased  so 
as  to  drive  off  the  remaining  water.  With  care  this  can  be  done  without 
breaking  the  tube,  when  it  will  be  seen  that  the  residue  is  a  nearly  pure 
white  powder.  If  the  test-tube  be  now  allowed  to  cool  and  some  water 
added,  the  powder  will  immediately  unite  with  a  part  of  it,  becoming 
blue  again.  On  exposure  to  the  air,  this  white  powder  takes  up  mois- 
ture enough  to  give  it  a  blue  color  in  a  very  short  time. 

Some  common  substances  contain  so  much  combined 
water  that  when  heated  they  appear  to  melt  and  assume 
the  liquid  form.  Ordinary  potassium  alum  and  sodium 
thiosulphate  (commonly  called  hyposulphite  of  soda)  show 
this  phenomenon. 

Ex.  13.  In  a  test-tube  carefully  heat  some  small  crystals  of  the 
sodium  thiosulphate.  It  will  be  seen  that  they  melt  very  readily  and  at 
a  low  temperature.  If  strongly  heated,  water  is  driven  off  and  can  be 
recognized.  The  liquid  obtained  by  melting  the  salt  in  its  water  of  crys- 
tallization behaves  as  a  supersaturated  solution,  which  can  be  shown  as 


GENERAL  CHEMISTRY.  11 

follows  :  After  liquefying  the  substance  close  the  tube  with  cotton  or  a 
cork,  and  stand  it  in  a  quiet  place  where  it  can  cool  down  without  any 
jar  or  agitation.  Under  these  conditions  the  substance  remains  as  a 
liquid.  If  now  the  stopper  be  removed  and  a  minute  crystal  be  dropped 
into  the  tube,  the  contents  solidify  immediately. 

Some  substances  holding  water  of  crystallization  can 
be  dehydrated  without  decomposition.  We  have  illustra- 
tions of  this  in  blue  vitriol,  borax,  crystal  soda,  alum, 
Glauber's  salt  and  others.  From  this  it  appears  that  the 
hydrated  and  anhydrous  forms  of  these  substances  are 
equally  stable.  But  other  substances  holding  water  are 
stable  only  in  this  condition  and  decompose  when  an 
attempt  is  made  to  separate  their  water.  Attention  will  be 
called  later  to  the  exact  chemical  composition  of  sub- 
stances crystallizing  with  water,  which  can  be  best  illus- 
trated by  means  of  formulas. 

Precipitation. 

It  has  been  shown  that  many  solid  substances  can  be 
dissolved  or  brought  into  solution  by  means  of  water  or 
other  liquid.  The  converse  of  this  will  now  be  illustrated. 
That  is,  it  will  be  shown  that  solutions  can  often  be  made 
to  give  up  their  dissolved  substance  by  other  methods 
than  by  evaporation  or  crystallization.  This  is  commonly 
accomplished  by  the  process  termed  precipitation.  By  this 
we  understand  a  process  in  which  a  dissolved  solid  or 
some  part  of  it  is  rendered  insoluble  and  settles  out  from, 
or  is  precipitated  from  the  solution.  Precipitates  are  usu- 
ally heavier  than  the  liquid  from  which  separated  and 
therefore  settle  to  the  bottom  of  the  containing  vessel.  A 
substance  in  solution  may  be  rendered  insoluble  in  several 
ways,  for  instance,  by  change  of  temperature,  by  adding 
to  the  solution  a  second  liquid  in  which  the  dissolved  sub- 
stance is  insoluble,  or,  most  commonly,  by  converting  the 
dissolved  substance  into  a  new  one,  insoluble  in  the  men- 
struum, by  addition  of  some  decomposing  reagent. 

As  an  illustration  of  precipitation  by  change  of  temper- 
ature the  following  test  may  be  made  : 

Ex.  14.     To  about  a  gram  of  calcium  tartrate  in  a  test-tube  add  10 


12  GENERAL  CHEMISTRY. 

cubic  centimeters  of  caustic  soda  solution.  The  solid  dissolves  by  shak- 
ing. When  a  clear  solution  is  obtained  boil  it  and  observe  that  a  gela- 
tinous precipitate  forms.  This  consists  of  the  calcium  tartrate,  ren- 
dered insoluble  by  increase  of  temperature. 

We  have  numerous  illustrations  of  precipitation  by  ad- 
dition of  a  second  liquid  to  the  solution,  and  the  methods 
are  frequently  applied  in  practical  analysis.  The  follow- 
ing experiments  will  serve  as  illustrations: 

Ex.  15.  Dissolve  some  crystallized  ferrous  sulphate  (common 
green  vitriol)  in  warm  water  and  filter  the  solution  to  make  it  perfectly 
clear.  To  some  of  this  clear  liquid  in  a  test-tube  add  an  equal  volume 
of  alcohol  and  observe  that  a  precipitate  of  the  sulphate  in  small  crystals 
settles  out.  A  more  satisfactory  result  can  be  obtained  by  pouring  the 
sulphate  solution  into  the  alcohol,  shaking  thoroughly  and  then  setting 
the  mixture  aside  several  hours. 

The  ferrous  sulphate  is  readily  soluble  in  water,  but  not 
in  alcohol;  hence  on  adding  the  latter  to  the  solution  the 
iron  compound  settles  out  unchanged.  Many  salts  may  be 
precipitated  from  their  aqueous  solutions  by  alcohol  in  the 
same  manner. 

Ex.  16.  By  aid  of  heat  dissolve  dextrin  or  other  gum  in  water.  To 
the  solution  add  some  alcohol  and  observe  the  precipitation  of  the  gum. 
Allow  this  to  subside,  which  may  require  hours;  pour  off  the  liquid  as 
far  as  possible  and  add  pure  water.  This  brings  the  gum  into  solution 
again,  indicating  that  the  addition  of  alcohol  had  rendered  it  only 
temporarily  insoluble. 

Ex.  17.  Dissolve  common  rosin  or  colophony  in  alcohol.  Pour 
some  of  the  clear  solution  into  a  test-tube  and  add  an  equal  volume  of 
water.  Precipitation  of  the  resin  substance  follows. 

Gums  are  soluble  in  water  usually,  but  not  in  alcohol. 
Resins  and  many  similar  bodies  are  soluble  in  alcohol  but 
not  in  water.  Hence  precipitation  takes  place  in  one  case 
by  adding  alcohol  to  the  aqueous  solution  and  in  the  other 
by  adding  water  to  the  alcohol  solution. 

In  the  illustrations  given  the  substance  precipitated 
separates  from  its  solution  in  practically  unchanged  con- 
dition. Precipitation  here  is  not  accompanied  by  decompo- 
sition. In  the  great  majority  of  cases,  however,  what  is 
termed  precipitation  depends  on  change  of  chemical 
composition,  and  is  brought  about  by  adding  to  a  solu- 


GENERAL  CHEMISTRY.  13 

tion  something,  usually  a  solution  of  another  substance, 
which  is  capable  of  producing  a  new  and  insoluble  body 
with  that  already  present.  The  new  body  formed  must 
therefore  settle  out  as  a  precipitate.  The  nature  of  this 
change  can  be  made  plain  best  by  a  few  simple  experi- 
ments. 

Ex.  18.  To  some  dilute  solution  of  blue  vitriol  (copper  sulphate) 
in  a  test-tube  add  an  equal  volume  of  solution  of  ammonium  sulphide 
and  shake.  From  the  mixture  of  the  blue  copper  solution  and  the 
nearly  colorless  or  yellow  sulphide  solution  we  obtain  a  black  substance 
which  is  evidently  not  the  original  copper  sulphate.  This  belief  is  con- 
firmed by  filtering  the  contents  of  the  test-tube.  A  yellow  or  brownish 
liquid  passes  through  the  paper  while  a  bulky  black  precipitate  remains. 
By  pouring  water  on  this  precipitate  it  fails  to  dissolve,  showing  its 
marked  difference  from  the  vitriol.  This  black  substance  is  known  as 
copper  sulphide,  and  in  many  important  properties  is  quite  unlike  the 
copper  sulphate  from  which  it  was  produced. 

Ex.  19.  To  a  solution  of  alum  in  a  test-tube  add  some  ammonia 
water  and  shake  the  mixture.  A  very  bulky  gelatinous  precipitate  forms 
and  gradually  settles  to  the  bottom  of  the  test-tube.  It  can  be  separated 
by  filtration  from  the  liquid  in  which  it  was  suspended,  and  when 
mixed  with  pure  water  is  found  to  be  insoluble.  In  appearance  and  in 
characteristic  properties  this  substance  is  very  different  from  the  origi- 
nal alum.  It  is  called  aluminum  hydroxide. 

Ex.  20.  By  the  aid  of  heat  dissolve  a  few  grams  of  powdered 
chalk  in  weak,  hydrochloric  acid  contained  in  a  test-tube  or  small 
beaker.  When  solution  is  complete  boil  a  few  minutes,  and  filter  if  the 
liquid  is  not  perfectly  clear.  We  have  now  a  solution  containing,  not 
the  chalk,  because  the  acid  decomposed  that,  but  calcium  chloride,  a 
new  substance.  That  we  have  here  a  substance  distinct  from  the  chalk 
can  be  shown  by  evaporating  some  of  the  solution  to  dryness  in  a  small 
porcelain  dish,  heated  on  a  sand  bath.  The  appearance  of  the  residue, 
and  the  fact  that  it  dissolves  in  water  while  the  chalk  does  not,  show  the 
distinct  nature  of  the  body.  Now,  to  the  remainder  of  the  solution  not 
evaporated  add  a  little  ammonia  water,  enough  to  make  it  impart  a  blue 
color  to  red  litmus  paper  after  stirring,  and  then  some  solution  of  am- 
monium carbonate.  This  will  produce  a  fine  white  precipitate  which 
settles  readily  to  the  bottom  of  the  vessel.  After  it  has  stood  some 
hours,  pour  off  the  liquid  as  far  as  possible,  and  collect  what  remains  on 
a  filter.  Allow  the  fine  white  precipitate  to  drain  thoroughly  and  then 
pour  water  over  it.  When  this  has  run  through  add  water  a  second 
time  and  wait  for  this  to  drain.  Then  stand  the  filter  aside  and  allow 
the  white  precipitate  to  become  thoroughly  dry,  which  may  require  sev- 
eral days.  On  examination  of  what  remains  it  will  be  found  to  have  the 
appearance  and  properties  of  the  original  chalk.  In  dissolving  the 
chalk  in  the  acid  it  was  observed  that  a  gas  was  given  off  and  this  gas  is 


14  GENERAL  CHEMISTRY. 

known  as  carbon  dioxide  or  carbonic  acid  gas.  One  element  of  the 
chalk,  however,  certainly  remained  behind,  because  a  solid  substance 
was  found  on  evaporating  the  solution.  It  appears,  therefore,  that  in 
the  formation  of  chalk  again  the  ammonium  carbonate  solution  must 
have  restored  in  precipitation  just  what  was  lost  when  the  chalk  went 
into  solution.  Pure  chalk  is  known  chemically  as  calcium  carbonate. 

An  important  peculiarity  of  precipitation  in  general  is 
shown  by  these  examples.  We  have  first  a  body  in  solu- 
tion with  its  particles  uniformly  distributed  among  those 
of  the  solvent.  A  condition  of  equilibrium  exists  of  such 
a  nature  that  any  tendency  of  the  particles  of  a  heavy 
body  to  sink  or  of  a  light  body  to  rise  and  float  is  exactly 
overcome.  Bodies  in  solution  resemble  gases  in  this 
respect,  that  their  particles  tend  to  separate  and  fill  all 
available  space  uniformly. 

We  have  seen  that  this  condition  of  equilibrium  maybe 
destroyed  in  several  ways — by  change  in  temperature,  by 
addition  of  a  new  liquid  in  which  the  dissolved  body  is  in- 
soluble, or  by  addition  of  a  certain  solution.  This  solution 
must  contain  a  substance  capable  of  forming  a  third  sub- 
stance insoluble  in  the  mixed  solutions.  We  may  have  an 
aqueous  solution  of  a  substance,  A,  and  a  second  aqueous 
solution  of  a  substance,  B,  but  it  does  not  follow  that  the 
product  of  the  action  of  A  on  B  should  also  be  soluble  in 
water.  It  often  happens  that  the  product  of  A  and  B  is 
insoluble.  For  example,  sodium  sulphate  and  barium 
chloride  are  both  easily  soluble  in  water,  but  on  mixing 
their  solutions  we  obtain  one  of  the  least  soluble  of  known 
substances,  barium  sulphate. 

The  formation  of  a  precipitate  in  a  liquid  is  not  an  in- 
stantaneous operation,  although  in  some  cases  the  interval 
between  the  addition  of  the  precipitant  and  the  formation  of 
a  precipitate  is  very  short.  The  precipitation  of  barium 
sulphate  is  an  illustration.  But  more  time  is  required  for 
the  completion  of  many  other  reactions,  as  will  be  seen  by 
the  following  experiment : 

Ex.  21.  Let  the  student  pour  some  dilute  solution  of  magnesium 
sulphate  into  each  of  three  tes't-tubes.  (This  solution  may  be  made  by 
dissolving  5  Gm.  of  the  crystallized  substance  in  100  Cc.  of  water.)  To 
the  first  test-tube  add  solution  of  barium  chloride;  a  precipitate  forms 
immediately,  apparently.  To  the  second  add  an  equal  volume  of  a 


GENERAL  CHEMISTRY.  15 

solution  of  calcium  chloride  containing  3  Gm.  in  100  Cc.  A  precipitate 
will  slowly  form.  To  the  third  test-tube  add  an  equal  volume  of  a  10 
per  cent  solution  of  ammonium  chloride,  then  some  ammonia  water, 
and  finally  a  few  drops  of  solution  of  sodium  phosphate.  In  time  a 
crystalline  precipitate  will  appear.  .  The  formation  of  this  precipitate 
may  be  aided  by  rubbing  the  sides  of  the  test  tube  with  a  glass  rod,  and 
the  insoluble  substance  settling  out  appears  first  in  the  form  of  minute 
glistening  specks,  which  grow  larger  and  finally  disclose  a  crystalline 
structure. 

It  is  evident  from  this  that  the  phenomenon  of  precipi- 
tation is  a  complex  one.  The  substance  we  recognize  as  a 
precipitate  is  not  immediately  formed,  but  is  a  growth,  the 
particles  we  see  being  formed  by  the  aggregation,  proba- 
bly, of  an  almost  infinite  number  of  smaller  particles.  The 
building  up  or  development  of  these  larger  particles  is 
often  greatly  aided  by  application  of  heat.  In  the  precip- 
itation of  barium  sulphate  in  the  above  experiment  the 
precipitate  remains  for  a  long  time  suspended  in  the  mixed 
liquid.  By  having  both  liquids  warm  it  settles  sooner, 
while  if  the  mixture  be  boiled  after  precipitation  the  white 
precipitate  will  settle  very  rapidly,  giving  evidence  of  the 
heaviness  and  compact  form  of  its  particles.  A  loose  pre- 
cipitate of  barium  sulphate,  as  it  is  produced  in  cold  solu- 
tions, cannot  be  easily  filtered.  The  particles  appear  to 
be  so  fine  that  they  can  pass  through  the  pores  of  ordi- 
nary filter  paper.  After  thorough  boiling,  however,  filtra- 
tion is  generally  easy,  the  particles  becoming  coarse 
enough  to  be  retained  on  the  filter. 

Ex.  22.  As  a  further  instructive  illustration  of  slow  precipitation 
the  following  experiment  may  be  made.  In  a  test-tube  mix  5  Cc.  of  a 
cold  solution  containing  about  6  Gm.  of  tartar  emetic  in  100  Cc.  with 
an  equal  volume  of  a  dilute  sodium  carbonate  solution  containing  about 
1  Gm.  in  100  Cc.  Apply  heat  to  the  mixture  and  observe  that  a  white 
precipitate  forms  immediately.  Now  repeat  the  experiment  using  the 
same  solutions  in  the  same  quantities,  but  have  both  as  cold  as  possible 
before  mixing  and  pour  the  soda  solution  into  the  other  very  slowly,  and 
with  little  agitation.  Close  the  test-tube  with  a  cork  and  leave  it  in  a 
quiet  place;  under  these  conditions  hours  may  elapse  before  the  slight- 
est trace  of  precipitation  appears.  On  shaking  the  tube,  pouring  out 
the  contents  or  slightly  warming,  a  precipitate  begins  to  form  and  soon 
becomes  heavy. 

The  behavior  here  recalls  that  already  observed  in  the 
experiments  with  supersaturated  solutions,  and  the  mixed 


16  GENERAL  CHEMISTRY. 

liquid  just  before  precipitation  was  in  a  supersaturated 
condition;  the  subsidence  of  the  precipitate  relieves  this. 
In  the  great  majority  of  cases  of  precipitation  the  time  in 
which  supersaturation  can  be  said  to  exist  is  extremely 
short,  so  as  to  escape  observation. 

Precipitates  are  distinguished  from  each  other  by  color, 
size  of  particles,  apparent  density  of  particles,  rapidity  of 
formation  and  subsidence,  degree  of  insolubility  and  in 
many  other  ways.  No  two  precipitates  are  exactly  alike 
and  we  have  therefore  in  the  phenomenon  of  precipitation 
something  of  value  for  the  recognition  of  substances,  In 
analytical  chemistry  precipitation  plays  a  very  important 
part  as  a  means  of  separation  and  identification.  In  chem- 
ical industry  many  substances  are  secured  by  precipitation 
from  solutions  containing  them.  In  the  pages  to  follow 
these  and  other  applications  will  be  abundantly  illustrated. 

Distillation. 

It  was  explained  in  the  beginning  of  this  chapter  that 
many  substances  are  capable  of  existing  in  three  forms,  as 
solids,  liquids  and  gases,  or  vapors.  The  conversion  of  a 
solid  or  liquid  into  a  vapor  is  usually  termed  vaporization 
and  may  take  place  spontaneously,  or,  commonly,  by  the 
application  of  heat. 

The  operations  of  vaporization  and  subsequent  conden- 
sation of  the  vapor  to  the  liquid  or  solid  condition  again, 
taken  together,  constitute  what  is  termed  distillation.  When 
sufficient  heat  is  applied  to  water  in  a  flask,  it  boils  and 
steam  is  formed  which  escapes  from  the  mouth  of  the  flask. 
The  production  and  escape  of  the  steam  alone  do  not  con- 
stitute distillation,  but  if  the  neck  of  the  flask  is  closed 
with  a  perforated  cork,  or  rubber  stopper,  through  which  a 
long  glass  tube,  bent  downward  after  leaving  the  flask, 
passes,  some  or  perhaps  nearly  all  of  the  steam  will  con- 
dense to  form  water  which  may  be  collected  from  the  end 
of  the  tube. 

The  flask  and  bent  glass  tube  constitute  a  rude  distil- 
ling apparatus  which  can  be  readily  constructed  by  the 
student  and  used  for  the  following  experiment.  See  Fig.  1. 


GENERAL  CHEMISTRY. 


17 


Ex.  23.  Into  the  neck  of  a  glass  flask  holding  about  300  Cc.  fit  a 
good  cork  or  rubber  stopper  having  a  perforation  at  least  three-eighths 
of  an  inch  in  diameter.  Next  select  a  piece  of  glass  tubing  just  wide 
enough  to  fit  the  hole  in  the  stopper  snugly  and  about  three  feet  long, 
and  melt  the  rough  ends  in  the  flame  of  the  Bunsen  burner  to  remove 
the  sharp  edges.  Then  about  four  inches  from  one  end  of  the  tube  make 
a  bend  by  heating  it  in  a  broad  flame  until  it  is  soft  enough  to  be  bent 
so  that  the  shorter  limb  makes  an  angle  of  about  60°  with  the  longer. 
(The  method  of  working  glass  tubing  must  be  learned  from  the 
instructor.)  This  shorter  limb  passes  through  the  perforation  in  the 
stopper.  Pour  into  the  flask  about  150  Cc.  of  water,  add  some  salt, 
enough  to  give  a  strong  taste,  and  then  a  little  indigo  solution  or  other 
highly  colored  liquid.  Insert  the  stopper  with  its  bent  glass  tube,  sup- 
port the  flask  on  a  sand-bath  or  wire  gauze  by  means  of  iron  rings  or 


FIG.    1. 


clamps  and  then  apply  heat  slowly,  below  the  sand  or  gauze.  The  water 
in  the  flask  becomes  hot  and  finally  begins  to  boil.  Steam  passes  up  into 
the  bent  tube  and  then  condenses  readily,  if  the  heat  applied  is  not  too 
strong. 

Allow  a  few  drops  of  the  condensed  liquid  to  fall  from  the  end  of  the 
tube  and  then  collect  what  follows  in  several  perfectly  clean  test-tubes. 
It  will  be  observed  that  the  colored  liquid  in  the  flask  yields  a  colorless 
distillate,  and  also  that  the  latter  is  free  from  salty  taste.  To  the  water 
collected  in  one  of  the  tubes  add  a  few  drops  of  solution  of  silver  nitrate. 
No  change  should  follow.  To  some  water  containing  salt,  as  poured  into 
the  flask,  add  silver  nitrate  and  observe  that  a  heavy  white  curdy  pre- 
cipitate forms.  These  experiments  show  that  the  salt  which  gives  the 
characteristic  taste  and  the  white  precipitate  with  the  silver  solution 
does  not  pass  over  with  the  steam.  It  is  not  readily  volatile.  The  sub- 
stance of  the  colored  liquid  is  likewise  nonvolatile. 


18  GENERAL  CHEMISTRY 

This  experiment  illustrates  the  manner  of  separation  of 
a  volatile  from  a  nonvolatile  substance  in  general.  Water 
and  other  liquids  are  commonly  purified  by  distillation; 
that  is,  they  are  in  this  manner  separated  from  solid  sub- 
stances they  hold  in  solution.  Practically,  the  very  simple 
distillation  apparatus  used  in  the  experiment  cannot  often 
be  employed.  In  most  cases  the  condensation  of  the 
vapor  would  be  quite  imperfect.  Instead  of  the  simple 
glass  tube  a  more  elaborate  condenser  is  usually  attached 
to  the  flask  or  still,  and  the  forms  best  known  should  be 


FIG.  2. 

on  exhibition  in  the  laboratory.  In  the  great  majority  of 
cases  in  practice  condensation  is  effected  by  passing  the 
vapor  through  a  straight  or  worm  tube  surrounded  by 
flowing  cold  water.  Forms  of  distilling  apparatus  are 
shown  in  the  figures  2  and  3. 

In  practical  laboratory  work  the  operation  of  distillation 
is  a  very  common  one.  By  it  liquids  may  often  be 
separated  from  solids,  volatile  solids  from  nonvolatile, 
easily  volatile  liquids  from  such  as  are  not  readily  vapor- 
ized, or  solids  volatile  at  a  low  temperature  from  those 
volatile  at  a  high  temperature.  In  these  latter  operations 


GENERAL  CHEMISTRY. 


19 


the  method  known  as  fractional  distillation  is  often  em- 
ployed. The  principal  applications  of  fractional  distilla- 
tion are  in  organic  chemistry,  but  a  simple  illustration  may 
be  given  here. 

An  approximate  separation  of  water,  alcohol  and  ether 
may  readily  be  made,  because  these  substances  boil  at 
very  different  temperatures.  Ether  boils  at  35°  C.,  alcohol 
at  78.5°  C.  and  water  at  100°  C.  Therefore  if  a  mixture  of 
these  substances  be  distilled  from  a  flask  and  the  distillate 
collected  in  small  portions  or  fractions  it  is  evident  that 


FIG.    3. 


the  first  fractions  will  consist  mainly  of  ether  and  the  last 
of  nearly  pure  water,  while  in  the  fractions  collected  near 
80°  C.  the  alcohol  will  be  in  excess.  A  sharp  separation 
is  not  possible,  because  the  ether,  beginning  to  boil  at  35°, 
carries  with  it  in  the  form  of  vapor  some  alcohol  and  even 
a  little  water.  In  turn  the  vapor  of  alcohol  carries  with  it 
some  water  vapor,  so  that  the  fractions  are  far  from  pure  at 
first. 

A  familiar  illustration  of  the  application   of  fractional 
distillation  on  the  large   scale  is  found  in  the  refining  of 


20  GENERAL  CHEMISTRY. 

crude  petroleur.i,  which  consists  of  a  mixture  of  many 
liquids  of  different  boiling  points.  A  full  explanation  of 
the  phenomenon  of  fractional  distillation  would  be  out  of 
place  at  this  time,  but  may  be  found  in  the  larger  works 
on  organic  chemistry,  a  subject  for  later  study. 


Chemical  and  Physical  Changes. 

In  our  experiments  on  the  precipitation  of  barium  sul- 
phate and  several  other  substances  we  had  an  illustration 
of  what  is  termed  a  chemical  change.  In  the  melting  of 
iodine  or  sulphur  we  had  an  illustration  of  a  physical 
change.  The  distinction  between  chemical  and  physical 
changes  will  be  made  plain  by  a  few  simple  experiments. 

Ex.  24.  Heat  some  pieces  of  bright  copper  or  iron  wire  in  the  hot 
flame  of  the  Bunsen  burner.  Observe  that  the  surfaces  become  tar- 
nished and  that  by  repeating  the  operation  several  times  a  dark,  brittle 
scale  is  formed  which  can  be  easily  rubbed  or  scraped  off  with  a  knife. 
Next  heat  a  small  piece  of  magnesium  wire  in  the  Bunsen  flame.  When 
it  becomes  quite  hot  it  burns  with  a  white,  dazzling  light,  giving  off  a 
white,  cloud-like  substance  which  finally  settles  down  as  a  powder. 
Finally,  heat  a  piece  of  platinum  wire  in  the  same  hot  flame.  It  will  be 
seen  that  while  held  above  the  lamp  the  metal  becomes  very  hot  and 
bright  red,  but  no  evidence  of  scaling  or  formation  of  fumes  is  seen. 
On  removing  the  wire  from  the  heat  it  resumes  its  former  color,  and  as 
far  as  can  be  seen  it  is  in  no  manner  different  from  what  it  was  before 
heating. 

These  simple  experiments  are  very  instructive.  By 
heating  the  copper  or  iron  it  was  evident  that  something 
new  was  formed  with  properties  different  from  those  of  the 
original  metal.  The  black  scale  scraped  from  the  iron  or 
copper  is  brittle  and  hard,  while  the  metals  are  ductile. 
The  white  powder  formed  from  the  magnesium  is  evi- 
dently quite  distinct  from  the  metal  and  it  becomes  appar- 
ent that  in  the  operation  of  heating  something  has  been 
lost  by  the  metals  or  absorbed  by  them  which  changes 
them  into  new  substances.  In  an  early  period  of  chemical 
study  it  was  held  that  under  the  influence  of  heat  metals 
lost  something.  It  is  now  known  that  instead  of  losing 
weight  the  copper,  the  iron  and  the  magnesium  take  up 
something  from  the  air  which  converts  them  into  new  sub- 


GENERAL  CHEMISTRY.  21 

stances,  with  an  increase  instead  of  a  loss  in  weight.  This 
absorption  of  something  from  the  air  with  increase  in 
weight  constitutes  a  radical  change  in  the  substance  under 
experimentation,  a  change  in  which  its  characteristic  prop- 
erties disappear,  giving  place  to  equally  marked  properties 
in  the  new  substance.  The  identity  of  magnesium  is  so 
completely  lost  in  the  white  powder  formed  by  burning 
that  the  recognition  of  the  relation  of  the  two  substances 
is  regarded  as  one  of  the  triumphs  of  early  chemical 
investigation.  Changes  as  far  reaching  as  these,  changes 
involving  frequently  a  loss  of  identity,  are  spoken  of  as 
cJiemical  changes. 

Even  superficial  examination  shows  that  no  radical  al- 
teration takes  place  in  the  platinum  during  the  heating 
operation.  The  change  there  was  merely  a  temporary  one, 
involving  no  real  loss  of  identity.  Such  changes  are  termed 
physical. 

Ex.  25.  In  a  test-tube  mix  some  flowers  of  sulphur  with  fine  cop- 
per turnings.  Gradually  apply  heat  to  the  mixture.  At  first  the  sul- 
phur melts  and  becomes  very  dark  colored.  As  the  temperature  grows 
higher  a  point  is  reached  where  a  combination  suddenly  takes  place 
between  the  sulphur  and  the  copper  which  is  shown  by  the  glowing  of 
the  latter.  The  copper  seems  to  burn  in  the  atmosphere  of  sulphur  in 
the  tube.  After  this  experiment  the  tube  is  allowed  to  cool  and  may  be 
broken.  In  place  of  the  bright,  ductile  copper,  a  black,  brittle  body  is 
found,  which  evidently  has  but  few  of  the  properties  of  the  original 
metal.  The  substance  formed  here  is  termed  copper  sulphide,  and  in 
its  production  we  have  a  typical  chemical  change. 

Ex.  26.  Pour  some  solution  of  blue  vitriol  into  a  beaker,  and  add 
to  it  a  little  dilute  sulphuric  acid.  Next  add  a  few  small  fragments  of 
granulated  zinc,  and  allow  the  beaker  to  stand  half  an  hour.  It  will 
soon  be  recognized  that  a  change  is  going  on  in  the  beaker,  as  the  zinc 
becomes  coated  with  a  dark,  spongy  mass,  in  color  suggesting  the  cop- 
per. It  will  be  observed,  also,  that  the  blue  color  of  the  liquid  gradu- 
ally becomes  fainter,  and  finally  that  it  may  disappear  entirely.  (This 
depends  on  the  amount  of  zinc  taken.)  By  removing  the  spongy  mass 
from  the  beaker,  washing  and  drying  it,  the  properties  of  copper  maybe 
recognized.  Meanwhile  it  should  be  observed  that  the  zinc  has  wholly 
or  in  part  disappeared 

We  have  here  a  very  curious  chemical  change,  which 
will  find  a  fuller  explanation  later.  But  it  may  be  said  now 
that  the  zinc  appears  to  go  into  the  solution,  while  the 


22  GENERAL  CHEMISTRY 

copper  of  the  blue  vitriol  or  copper  sulphate  solution  is 
precipitated.  At  any  rate,  the  identity  of  the  solution  is 
destroyed  with  the  loss  qf  its  copper. 

Ex.  27.  Over  some  small  nails  or  tacks  in  a  beaker  pour  some 
dilute  sulphuric  acid.  Very  soon  an  evolution  of  gas  is  observed,  and 
after  a  time  the  metal  will  have  disappeared.  A  light  green  liquid 
results,  and  this  evidently  contains  the  iron  in  a  dissolved  form.  We 
have,  in  fact,  a  solution  of  green  vitriol  or  ferrous  sulphate,  which 
could  be  separated  by  crystallization. 

It  will  not  be  necessary  to  multiply  instances,  as  enough 
has  been  given  to  show  what  is  characteristic  in  so- 
called  chemical  changes.  Iron,  under  the  action  of  a  high 
heat  or  by  treatment  with  an  acid,  becomes  changed  mate- 
rially by  conversion  into  something  else  which  is  not  iron, 
but  which  contains  iron.  Under  the  action  of  a  strong 
magnet  iron  becomes  likewise  changed,  but,  as  we  know, 
only  temporarily.  On  the  removal  of  the  magnet  the  iron 
assumes  its  original  nature  and  important  properties^  and 
gives  little  or  no  evidence  of  the  physical  change  through 
which  it  has  passed. 

A  careful  study  of  the  common  chemical  changes  or 
reactions  shows  that  we  can  make  three  general  divisions  of 
them.  We  have  first,  reactions  of  decomposition  in  which 
a  single  substance  is  broken  up  or  decomposed  so  as  to 
yield  two  or  more  other  substances.  We  have  many  illus- 
trations of  this.  For  example,  in  "burning"  lime  the 
common  rock  known  as  limestone  is  strongly  heated  until 
it  becomes  decomposed,  yielding  a  residue  called  quicklime 
which  on  addition  of  water  becomes  slaked  lime,  and  a 
gas  called  carbonic  acid  gas  or  carbon  dioxide.  This  is 
generally  allowed  to  escape.  The  action  of  the  heat  here 
is  to  effect  disintegration,  but  it  adds  nothing  in  the  form 
of  matter  to  the  limestone.  A  simple  experiment  will  be 
given  in  which  the  decomposition  of  a  substance  is  easily 
shown. 

Ex.  28.  In  a  small  test-tube  heat  a  few  grams  of  r^d  mercuric 
oxide  to  a  high  temperature  The  substance  darkens  and  finally  begins 
to  break  up,  as  may  be  readily  shown  by  two  phenomena.  In  the  test- 
tube  above  the  heated  powder  a  deposit  of  fine  metallic  globules  col- 
lects and  this  is  easily  recognized  as  mercury  itself.  If  while  a  strong 


GENERAL  CHEMISTRY.  23 

heat  is  being  applied  a  glowing  splinter  be  held  just  within  the  mouth 
of  the  test-tube  it  will  burst  into  flame  and  burn  with  great  brilliancy. 
This  shows  that  in  addition  to  the  metallic  globules  furnished  by  the 
red  compound  a  gas  is  liberated,  for  only  a  gas  could  exhibit  the  behav- 
ior just  mentioned.  This  experiment  will  be  taken  up  again. 

We  have  here  a  characteristic  reaction  of  decomposi- 
tion without  the  aid  of  outside  matter  in  which  we  obtain 
from  a  heavy  red  powder  a  silvery  liquid,  mercury,  and  a 
gas,  oxygen.  Such  reactions  are  frequently  termed  analyt- 
ical reactions  because  they  consist  in  an  analysis  or  break- 
ing up  of  something. 

We  have  next  reactions  of  just  the  opposite  character, 
that  is,  reactions  in  which  two  or  more  substances  com- 
bine to  form  a  third  body.  Our  experiment  on  heating 
the  copper  and  sulphur  is  an  illustration  of  reactions  in 
this  group.  At  a  high  temperature  the  two  substances 
were  shown  to  combine,  forming  a  new  compound  called 
copper  sulphide  or  sulphide  of  copper.  A  second  illustra- 
tion is  furnished  by  the  rusting  of  iron.  Here  the  metal 
combines  with  something  from  the  air  (oxygen),  produc- 
ing oxide  of  iron.  It  should  also  be  mentioned  that  under 
certain  conditions  the  reactions  described  above  by  which 
limestone  and  the  red  oxide  of  mercury  were  each  sepa- 
rated into  two  substances  are  reversible.  That  is,  the 
lime  and  carbonic  acid  gas  may  be  combined  to  form  lime- 
stone, and  metallic  mercury  and  oxygen  to  form  the  red 
oxide.  Many  such  reactions  are  known  and  they  are  some- 
times called  synthetical  reactions. 

The  most  important  and  numerous  of  our  chemical 
changes  belong  to  a  third  group,  however.  Here  two  or 
more  substances  react  on  each  other  to  produce  two  or 
more  new  substances.  Several  illustrations  of  such  de- 
compositions were  given  above  under  the  head  of  precipi- 
tation. Another  may  be  given  here. 

Ex.  29.  Take  a  few  grams  of  granulated  zinc  in  a  beaker  and  pour 
over  it  some  dilute  sulphuric  acid.  An  effervescence  begins  immedi- 
ately showing  the  escape  of  .a  gas,  in  some  manner  produced  by  the 
action  of  the  zinc  on  the  acid.  This  is  an  evidence  of  the  formation  of 
at  least  one  new  substance,  because  the  gas  can  be  neither  the  zinc  nor 
the  acid.  It  will  be  readily  seen  that  the  acid  dissolves  the  zinc,  that 
is,  that  a  solution  is  formed,  and  when  the  action  is  complete,  which  is 


24  GENERAL  CHEMISTRY. 

shown  by  the  disappearance  of  the  metal,  pour  some  of  the  solution  into 
a  small  porcelain  evaporating  dish  on  a  sand-bath,  and  apply  heat  to 
drive  off  everything  volatile.  Finish  the  concentration  in  a  fume 
closet,  applying  finally  a  strong  heat.  A  white  residue  will  be  left 
which  is  plainly  neither  zinc  nor  sulphuric  acid.  We  have,  therefore, 
in  this  case  the  formation  of  a  gas  and  a  white  solid  substance  from  a 
metal  and  an  acid  liquid,  which  can  readily  be  shown  to  be  wholly 
volatile.  The  gas  is  hydrogen  and  the  white  solid  is  zinc  sulphate. 

Ex.  30.  In  a  test-tube  take  about  10  Cc.  of  strong  "  sugar  of  lead" 
solution  (solution  of  lead  acetate).  Heat  to  boiling  and  observe  the 
odor.  In  another  test-tube  take  an  equal  volume  of  dilute  sulphuric 
acid,  boil  and  observe  the  odor.  Mix  the  hot  liquids.  A  white  precipi- 
tate forms  which  certainly  does  not  resemble  either  one  of  the  original 
substances.  It  will  be  noticed  also  that  the  mixed  liquid  emits  a  strong 
odor  of  vinegar  or  acetic  acid.  To  show  more  clearly  what  has  hap- 
pened allow  the  precipitate  to  settle  in  the  test-tube  and  then  pour  the 
liquid  above  it  through  a  filter.  Heat  the  filtered  liquid  to  the  boiling 
point  and  observe  that  the  odor  is  very  strong  and  characteristic.  Then 
add  water  to  the  residue  in  the  test-tube,  warm  and  pour  the  mixture  on 
the  same  filter,  and  wash  it  several  times  by  pouring  on  water.  That 
this  residue  is  not  the  sulphuric  acid  is  evident,  that  it  is  not  the  lead 
acetate  is  shown  by  the  fact  that  it  is  not  soluble  in  the  water  poured 
over  it,  while  the  lead  acetate  is,  readily.  We  have,  therefore,  in  this 
case  the  production  of  an  insoluble  residue  and  a  volatile  liquid  suggest- 
ing vinegar.  The  residue  is  known  as  lead  sulphate  and  the  volatile 
liquid  is  acetic  acid. 

It  can  be  readily  shown  that  in  many  common  precipita- 
tions two  substances  give  rise  to  two  new  ones,  but  these 
illustrations  are  sufficient  for  the  purpose  at  present. 

The  fact  that  we  obtained  above  two  bodies  from  one 
by  application  of  heat  is  sufficient  proof  of  the  compound 
nature  of  that  body. 

It  is  plain  that  a  body  formed  by  the  union  of  two  must 
be  compound,  containing  at  least  two  component  parts. 
Finally,  when  we  obtain  two  new  substances  by  the  action 
on  each  other  of  two  different  bodies  it  is  evident  that  one 
of  them  at  least  must  be  compound. 

These  considerations  can  be  illustrated  by  symbols  as 
follows  : 

Let  AB  represent  a  compound  body  which  under  cer- 
tain conditions  is  broken  up  into  its  component  parts,  A 
and  B.  Then  we  can  write 

AB  yields  A  -f  B. 


GENERAL  CHEMISTRY.  25 

In  the  second  case  we  have  the  reverse  of  this  reaction, 
that  is 

A  +  B  yields  AB. 

In  the  third  case  we  go  further  and  have  evidently  more 
component  parts  than  A  and  B  to  consider.  We  have 
evidently  C  and  D  also,  and  we  can  express  our  result  in  a 
general  way  as  follows  :  The  bodies  AB  and  CD  act  on 
each  other  and  make  AD  and  BC,  or 

AB  +  CD  yields  AD  +  BC. 

At  present  no  reason  appears  why  we  should   not  write 
instead  of  the  above  this: 

AB  +  CD  yields  AC  -f  BD. 

But  later  a  meaning  will  be  attached  to  these  symbols 
which  will  render  plain  just  what  does  take  place.  In  cer- 
tain cases  we  can  express  our  reaction  in  this  manner: 

A  +  BC  yields  AB  +  C. 

In  this  instance  only  one  of  the  bodies  entering  the 
reaction  is  considered  as  a  compound  one. 

This  is  BC,  which  the  simple  substance,  A,  decomposes 
into  the  new  compound  body,  AB,  and  the  new  simple 
body,  C. 

Chemists  usually  represent  these  changes  by  what  are 
termed  equations,  as, 

A  +  B  =  AB 

AB  =  A  -f  B 
AB  -h  CD  =  AD  +  BC 

B  +  CD  =  D  +  CB. 

What  is  written  to  the  left  of  the  =  sign  represents 
that  which  is  taken,  and  the  result  of  the  chemical  change 
is  shown  on  the  right  hand  side  of  the  sign. 

In  these  equations  the  letters  A,  B,  C  and  D  represent 
the  elements  or  parts  of  compounds  which  take  part  in  the 
reactions.  Their  full  meaning  will  be  explained  later. 

Conditions  of  Chemical  Change. 

The  conditions  under  which  chemical  changes  take 
place  are  different  in  different  combinations.  In  some  of 


20  GENERAL  CHEMISTRY. 

the  illustrations  given  above  it  has  been  shown  that  cer- 
tain substances  can  be  made  to  combine  by  the  aid  of 
heat,  while  in  other  cases,  decomposition  is  effected  by 
heat.  These  were  cases,  however,  in  which  dry  substances 
were  taken  for  experiment.  At  the  ordinary  temperature 
such  bodies  enter  into  combination  or  decompose  as  a  rule, 
but  slowly.  A  few  illustrations  will  be  given  in  which 
solid  bodies  are  combined  by  friction. 

Dry  Reactions.  The  following  three  experiments  are 
simple  cases  : 

Ex.  31.  Rub  together,  in  a  mortar,  about  equal  weights  of  corro- 
sive sublimate  and  potassium  iodide.  A  bright  red  compound  results, 
which  is  different  from  the  substances  giving  rise  to  it,  not  only  in  color 
but  in  solubility,  as  may  be  shown.  Add  some  water  to  the  contents  of 
the  mortar  and  stir  well.  A  red  precipitate  remains.  This  is  a  new 
compound,  mercuric  iodide. 

Ex.  32.  Rub  together  in  a  mortar,  minute  quantities  (a  few  milli- 
grams of  each  only}  of  sugar  and  potassium  chlorate.  The  substances 
react  on  each  other  violently,  producing  an  explosion.  If  large  quanti- 
ties were  used  the  experiment  would  be  very  dangerous.  The  chemical 
change  taking  place  here  results  in  the  formation  of  bodies  very  differ- 
ent from  the  sugar  or  the  potassium  chlorate. 

Ex.  33.  Mix  in  a  mortar,  by  means  of  a  piece  of  paper,  or  card- 
board, about  equal  weights  of  dry  slaked  lime  and  ammonium  chloride. 
At  first  no  change  should  be  noticed,  but  on  applying  some  pressure  in 
mixing,  as  when  the  two  substances  are  ground  together  with  a  pestle,  a 
change  rapidly  takes  place  in  which  ammonia  is  liberated,  as  shown  by 
the  smell.  The  mass  becomes  moist  as  the  rubbing  is  continued.  The 
nature  of  this  reaction  will  be  explained  later. 

Reactions  in  the  dry  way,  as  illustrated  above,  are  in- 
teresting but  not  very  common.  A  few  have  practical 
importance,  but  by  far  the  greater  number  of  chemical 
changes  with  which  we  are  acquainted,  take  place  in  solu- 
tion. 

Reactions  in  Solution.  Our  experiments  on  precipi- 
tation are  illustrations  of  these,  but  others  may  be  given. 

Ex.  34.  Dissolve  very  small  amounts,  as  in  Ex.  31,  of  mercuric 
chloride  (corrosive  sublimate)  and  potassium  iodide  in  water  and  mix 
the  solutions.  The  deep  red  precipitate  results  immediately. 


GENERAL  CHEMISTRY.  27 

Ex.  35.  Mix  the  slaked  lime  and  ammonium  chloride  mentioned 
in  Ex.  33"  with  water,  and  warm  gently.  The  strong  ammonia  odor 
soon  appears.  Vary  this  experiment  then  by  using  instead  of  the  lime, 
solutions  of  caustic  soda  and  sodium  carbonate,  which  likewise  liberate 
the  ammonia.  In  these  experiments  the  ammonium  chloride  is  com- 
pletely decomposed,  the  volatile  ammonia  escaping. 

Ex.  36.  Mix  together  on  a  piece  of  dry  paper  some  sodium  bicar- 
bonate ("  baking  soda  ")  and  some  dry  powdered  tartaric  acid.  As  long 
as  the  mixture  is  kept  perfectly  dry  no  apparent  change  takes  place. 
The  substances  do  not  seem  to  react  on  each  other,  and  in  fact  the  mix- 
ture may  be  kept  dry  almost  indefinitely.  But  if  it  is  thrown  into  a 
beaker  and  water  added  a  lively  effervescence  begins,  due  to  the  escape 
of  carbonic  acid  gas  from  the  decomposed  bicarbonate.  The  addition 
of  water  brings  both  substances  taken  into  solution,  in  which  condition 
they  act  readily  on  each  other. 

The  above  is  a  typical  experiment,  as  many  changes 
take  place  in  the  same  manner.  The  action  of  the  common 
baking  powders  depends  on  the  behavior  here  illustrated. 
A  large  number  of  substances  seem  to  have  no  action  on 
each  other  when  mixed  in  the  perfectly  dry  condition,  but 
when  dissolved  mutual  decomposition  begins.  In  the 
above  experiment  it  is  not  merely  the  soda  which  is  altered, 
as  shown  by  the  escape  of  gas,  but  the  tartaric  acid  suffers 
a  change  also.  It  is  converted  into  a  neutral  body,  that  is, 
one  without  acid  properties. 

It  seems  to  be  true  that  in  solution  the  particles  of  dis- 
solved substance  are  brought  into  a  condition  in  which 
tkey  move  with  great  freedom  and  may  thus  be  brought 
into  intimate  contact  with  each  other,  which  is  not  the  case 
as  long  as  they  are  in  the  dry  form.  In  general,  solutipn 
is  favorable  to  chemical  change,  and  we  therefore,  as  far 
as  possible,  dissolve  the  substances  we  wish  to  combine 
with  each  other,  for  the  production  of  new  substances. 

Reactions  of  Gases.  Not  only  have  we  reactions 
between  solids  and  reactions  between  liquids,  but  we  have 
also  some  well  marked  cases  of  reactions  between  gases. 
A  few  of  these  have  practical  importance.  The  following 
experiment  will  serve  as  an  illustration: 

Ex.  37.  By  means  of  a  glass  rod  place  a  drop  of  strong  hydro- 
chloric acid  on  one  side  of  the  bottom  of  a  dry  beaker.  Clean  and  dry 


28  GENERAL  CHEMISTRY. 

the  rod,  and  with  it  put  a  drop  of  strong  ammonia  water  on  the  opposite 
side  of  the  beaker.  The  first  drop  contains  hydrochloric  acid  gas,  tne 
second  ammonia  gas.  Some  of  each  gas  leaves  the  liquid  in  which  it  is 
dissolved  and  the  two  unite  in  the  beaker,  producing  white  fumes.  These 
white  fumes  consist  of  ammonium  chloride,  a  solid  substance,  which  is 
finally  deposited  on  the  walls  of  the  beaker.  After  placing  the  two 
drops  in  the  beaker  it  should  be  covered  with  a  glass  plate. 

Several  other  gases  combine  readily  in  the  same  man- 
ner. In  some  cases  the  products  formed  are  also  gases; 
in  other  cases  they  are  liquids,  while  sometimes,  as  in 
the  above  experiment,  they  are  solids.  In  a  few  cases 
the  combination  takes  place  readily  and  spontaneously,  but 
in  other  cases  it  must  be  brought  about  by  special  means. 
A  mixture  of  hydrogen  gas  with  oxygen  gas  may  be  kept 
at  the  ordinary  temperature,  but  by  application  of  heat 
and  by  other  means  the  two  gases  combine  with  explosive 
violence,  if  in  certain  proportions,  forming  water.  A 
mixture  of  hydrogen  and  chlorine  gases  may  be  kept  in 
the  dark,  but"  if  brought  into  the  sunlight  a  sudden  com- 
bination or  explosion  follows,  hydrochloric  acid  gas  being 
formed.  In  later  chapters  other  illustrations  of  gaseous 
combinations  will  be  given. 

In  all  cases  of  chemical  combinations,  whether  of  sol- 
ids, liquids  or  gases,  it  has  been  found  by  experiment  that 
the  substances  united  combine  in  certain  proportions  only. 
For  instance,  it  can  be  readily  shown  that  17  parts  of  am- 
monia gas,  by  weight,  combine  with  exactly  36.5  parts  by 
weight  of  hydrochloric  acid  gas. 

*  If  a  larger  amount  of  either  one  of  these  gases  were 
taken  with  the  given  weight  of  the  other,  this  excess  would 
fail  to  go  into  union  and  would  remain  in  the  free  state. 
In  the  reaction  between  hydrogen  and  oxygen  1  part  by 
weight  of  the  former  combines  with  8  parts  of  the  latter. 
In  combining  hydrogen  with  chlorine,  it  is  found  that  1 
part  by  weight  of  the  former  combines  always  with  35.5 
parts  of  the  latter,  but  not  with  more  or  less. 

In  Ex.  27,  it  was  shown  that  iron  is  dissolved  by 
sulphuric  acid.  By  careful  attention  to  details  it  can  be 
shown  that  the  amount  of  iron  which  can  be  dissolved  by 
a  given  weight  of  sulphuric  acid  is  absolutely  constant. 


GENERAL  CHEMISTRY.  29 

In  Ex.  29  zinc  is  dissolved  in  the  same  acid,  and  proper 
tests  show  that  the  weight  of  the  metal  dissolved  by  a 
given  weight  of  the  acid  is  constant  and  greater  than  the 
weight  of  iron  which  can  be  dissolved  in  the  same  acid. 

From  these  illustrations  it  would  appear,  therefore, 
that  in  our  chemical  combinations  we  have  quantitative  as 
well  as  qualitative  relations.  It  would  be  premature  to 
attempt  an  explanation  of  these  facts  now.  The  student 
should  bear  them  in  mind  and  look  for  an  explanation  in 
later  experiments. 


CHAPTER  II. 


OXYGEN,    HYDROGEN  AND  THEIR  COHPOUNDS. 

OXYGEN. 

WE  ARE  now  ready  to  begin  the  study  of  particular 
substances  somewhat  in  detail,  and  will  begin  with 
the  very  important  and  common  body  known   as  oxygen. 

Occurrence.  Oxygen  is  widely  distributed  throughout 
the  animal,  vegetable  and  mineral  kingdoms,  constituting 
about  one-half  of  the  total  weight  of  everything  we  are 
acquainted  with  in  and  above  the  earth's  crust.  It  makes  up 
eight  ninths  of  water  by  weight,  and  over  one-fifth  of  the 
atmosphere.  All  the  common  rocks  and  clay  contain  it  in 
combination,  while  in  such  common  substances  as  sugar, 
starch, the  fats,  albumin  and  woody  fiber  it  is  an  important 
constituent. 

History.  The  history  of  this  remarkable  body  is  inter- 
esting. While  now  easily  recognized  as  a  distinct  sub- 
stance it  must  be  remembered  that  the  earlier  chemists 
were  without  this  knowledge.  The  curious  properties 
which  will  be  shown  later  to  belong  to  oxygen  were  either 
overlooked  or  ascribed  to  something  else.  The  atmosphere, 
which  owes  its  most  important  properties  to  the  oxygen 
present,  was  supposed  to  be  a  simple  substance,  and  the 
common  phenomena  of  combustion  in  air  or  oxygen  were 
all  wrongly  interpreted.  However,  in  1774,  Priestley,  and, 
independently  of  him,  Scheele,  in  1775,  isolated  pure  oxy- 
gen from  compounds  containing  it  and  recognized  it  as  the 
important  element  of  the  air. 


GENERAL  CHEMISTRY.  31 

In  1781  Cavendish  announced  the  composition  of  water, 
and  a  little  later  the  great  French  chemist,  Lavoisier,  gave 
the  first  rational  explanation  of  the  behavior  of  oxygen  in 
combustion  and  respiration,  and  opened  the  way  for  the 
growth  of  the  science  of  modern  chemistry. 

Preparation.  Although  oxygen  is  abundantly  present 
in  well-known  materials  everywhere  obtainable,  we  secure 
it  in  the  pure  state  practipally  from  but  few  sources.  It  is 
obtained  from  the  air  at  very  slight  cost  by  processes 
which  cannot  be  explained  at  this  point,  but  only  where 
required  in  large  quantities  for  certain  manufacturing 
operations.  When  used  for  other  purposes  it  is  commonly 
made  by  the  decomposition  of  certain  compounds  contain- 
ing it — from  water,  from  the  red  oxide  of  mercury,  and  from 
potassium  chlorate,  for  instances. 

Two  experiments  will  be  here  given  to  illustrate  these 
operations: 

Ex.  38.  Repeat  Ex.  28  by  heating  a  small  amount  of  red  mercuric 
oxide  in  a  narrow  test-tube.  Observe  that  a  strong  heat  is  required  for 
the  decomposition  of  the  substance,  and  that  finally  a  glowing  splinter 
held  within  the  mouth  of  the  test-tube  rekindles  and  burns  brightly.  As 
already  intimated,  this  phenomenon  shows  that  a  gas  must  be  given  off 
by  the  action  of  heat  on  the  red  compound.  This  gas  is  oxygen,  and 
the  simple  experiment  illustrates  one  of  the  first  processes  given  for  its 
preparation. 

This  method  o1  liberating  oxygen  is  not  a  convenient 
one,  and  besides  is  very  expensive.  Experiment  shows 
that  216  parts  of  the  oxide  of  mercury  yield  only  16  parts 
of  oxygen,  and  a  high  temperature  is  required  to  separate 
this  from  the  mercury.  The  experiment  has  value  only  as 
an  illustration,  as  by  this  method  Priestley  first  secured 
the  gas  in  pure  condition. 

For  laboratory  uses  we  make  oxygen  generally  by  a 
process  indicated  by  the  following  experiment: 

Ex.  39.  In  a  test-tube  heat  about  10  Gm.  of  powdered  dry  potas- 
sium chlorate.  Move  the  test-tube  backward  and  forward  through  the 
flame,  turning  it  meanwhile  between  the  thumb  and  fingers  so  as  to 
avoid  cracking  it.  After  a  time  the  powder  melts  to  a  liquid,  which  by 
longer  application  of  heat  appears  to  boil.  Gas  bubbles  are  seen  to 
escape  from  it,  and  if  a  glowing  splinter  is  now  held  within  the  mouth 


32  GENERAL  CHEMISTRY. 

of  the  test-tube,  it  will  soon  burst  into  flame,  as  in  the  other  case  The 
high  heat  applied  here  decomposes  the  substance  taken,  and  oxygen  gas 
is  one  of  the  products.  What  remains  in  the  tube  will  appear  later. 

This  process  has  certain  drawbacks.  A  relatively  high 
temperature  is  required  for  the  breaking  up  of  the  chlorate 
and  the  reaction  is  somewhat  slow.  It  may  be  modified, 
however,  as  follows: 

Ex.  40.  Mix  about  equal  weights,  a  few  grams  of  each,  of  potas- 
sium chlorate  and  manganese  dioxide.  Heat  the  mixture  in  a  test-tube 
and  notice  that  the  liberation  of  gas,  which  kindles  the  flame  on  a  glow- 
ing coal,  begins  almost  immediately.  The  substance  does  not  melt,  but 
at  a  relatively  low  temperature  undergoes  decomposition  with  separation 
of  the  oxygen.  This  decomposition  is  very  rapid,  as  may  be  seen  by  the 
manner  in  which  splinters  of  wood,  once  ignited,  burn  at  the  mouth  of 
the  test-tube. 

The  last  experiment  shows   us  how  oxygen  may  be 
prepared  in  quantity,  which  will  now  be  tried. 

Ex.  41.  Mix  about  25  Gm.  of  dry  powdered  potassium  chlorate 
with  an  equal  weight  of  manganese  dioxide,  and  transfer  to  a  dry  and 
perfectly  clean  glass  flask,  holding  200  to  250  Cc.  By  means  of  a  per- 
forated stopper  connect  this  flask  with  a  bent  delivery  tube  arranged  as 
in  the  following  figure.  The  further  end  of  the  delivery  tube  is  bent 
upward  slightly  and  dips  beneath  the  surface  of  water  contained  in  a 
pneumatic  trough  or  earthenware  bowl.  The  trough  or  bowl  should  be 
nearly  filled  with  the  water.  The  flask  is  supported  on  a  sand-bath  by 
means  of  a  clamp  or  ring.  On  now  applying  heat  by  the  burner,  the 
mixture  in  the  flask  soon  becomes  hot  and  begins  to  decompose  as  shown 
by  the  escape  of  bubbles  of  gas  from  the  end  of  the  delivery  tube. 

Method  of  Collecting  the  Gas.  The  gas  is  most  readily  collected 
by  displacement  of  water,  and  the  directions  given  for  this  operation 
here  will  answer  for  many  later  experiments.  Have  at  hand  several 
wide  mouth  bottles,  holding  about  250-400  Cc.  each.  Fill  them  quite 
full  with  water  and  then  cover  them  with  squares  of  glass  in  such  a 
manner  as  to  exclude  all  air  bubbles.  When  this  is  done  each  bottle 
may  be  inverted  by  holding  it  with  one  hand,  while  the  plate  of  glass  is 
pressed  down  with  the  other.  The  mouth  of  the  bottle  in  this  position 
is  brought  beneath  the  surface  of  the  water  in  the  trough  and  the  plate 
removed.  The  bottle  remains  full,  its  contents  being  held  up  by  the 
atmospheric  pressure,  and  this  remains  true  in  whatever  position  the 
bottle  stands,  provided  its  mouth  is  always  beneath  the  surface  of  the 
trough  water.  It  may  therefore  be  held  just  over  the  end  of  the  deliv- 
ery tube  from  the  flask,  and  so  catch  the  gas  bubbles  as  they  ascend. 
Instead  of  holding  the  collecting  bottle  in  the  hand  it  is  much  better  to 
support  it  on  a  bridge  of  galvanized  iron,  which  has  a  perforation  in  its 
center  about  2  centimeters  in  diameter.  The  bubbles  can  pass  through 


GENERAL  CHEMISTRY. 


33 


this  opening  into  the  bottle  above.  When  the  bottle  is  full  of  gas  move 
it  to  one  side,  its  mouth  still  beneath  the  surface  of  the  water,  close 
with  the  square  of  glass  and  then  lift  it  out  of  the  water  and  stand  on 
the  table  in  an  upright  position.  If  the  plate  fits  the  bottle  well,  the 
gas  will  not  soon  escape.  Now  bring,  in  the  same  manner  other  bottles 
of  water  over  the  end  of  the  delivery  tube  and  collect  enough  gas  for  ah 
the  experiments  given  below.  The  first  bottle  of  gas  collected  may  be 
contaminated  by  air  from  the  generating  flask,  but  the  others  should 
contain  nearly  pure  oxygen.  On  completing  the  experiment  the  deliv- 
ery tube  should  be  taken  from  the  water  before  the  lamp  is  removed. 
Why? 


FIG.  4. 


With  the  bottles  of  oxygen  collected  as  just  explained, 
the  student  is  ready  to  make  some  simple  experiments  to 
illustrate  important  properties  of  the  gas. 

Potassium  chlorate  furnishes  39.2  per  cent  ot  its  weight 
of  oxygen  when  completely  decomposed.  At  the  ordinary 
temperature  1  Gm.  yields  about  300  cubic  centimeters  of 
gas.  For  the  preparation  of  quantities  of  the  gas  a  cop- 
per retort  is  commonly  employed,  and  from  this,  on  libera- 
tion, the  gas  is  led  into  a  large  reservoir  or  holder.  For 
certain  purposes  the  gas  requires  some  purification,  unless 
made  of  absolutely  pure  materials.  The  purification  can 
be  easily  effected  by  allowing  the  gas  on  leaving  the  gen- 
erating retort  to  bubble  through  a  wash  bottle  containing 


34  GENERAL  CHEMISTRY. 

a  solution  of  potassium  hydroxide.  The  arrangement  of 
generator,  wash  bottle  and  gas  holder  is  shown  in  Fig.  5. 

One  hundred  grams  of  the  chlorate  mixed  with  an 
equal  weight  of  manganese  dioxide  will  furnish  30  liters  of 
gas.  A  little  of  this  should  be  wasted,  however,  to  drive 
the  air  from  the  retort  and  wash  bottle  at  the  beginning  of 
the  operation. 

Before  attempting  to  heat  a  large  quantity  of  the  mixed 
chlorate  and  black  oxide,  as  explained  above,  a  trial  should 


FIG.  5. 

always  be  made  with  a  small  quantity  in  a  test-tube.  The 
black  oxide  has  been  occasionally  found  adulterated  or 
accidentally  mixed  with  charcoal  or  other  form  of  carbon, 
and  dangerous  explosions  have  resulted  from  the  careless 
use  of  such  a  product.  The  test-tube  experiment  would 
disclose  such  impurity  if  present. 

In  the  tests  made  in  experiments  39  and  40  it  was  dis- 
covered that  the  gas  supports  combustion  well,  as  shown 
by  the  rapid  and  brilliant  burning  of  the  splinter.  A  simi* 
lar  phenomenon  is  shown  in  the  next  test. 


GENERAL  CHEMISTRY.  35 

Ex.  42.  Attach  a  piece  of  charcoal  to  a  bent  wire,  bring  it  to  a 
glowing  condition  in  the  lamp  flame,  and  plunge  it  quickly  in  a  bottle  of 
the  gas.  It  burns  brilliantly,  throwing  off  showers  of  sparks.  Pour 
some  lime-water  in  the  bottle  at  the  end  of  the  reaction,  close  with  a 
glass  plate  and  shake.  A  white  precipitate  is  formed,  owing  to  the  com- 
bination of  a  constituent  of  the  lime-water  (calcium  hydroxide),  with 
the  gas  produced  by  the  combustion  of  the  charcoal.  The  gas  is  carbon 
dioxide.  The  precipitate  consists  of  calcium  carbonate. 

This  experiment  gives  us  a  typical  example  of  what  is 
termed  combustion.  The  oxygen  in  the  bottle  goes  into 
intimate  union  or  chemical  combination  with  the  hot  char- 
coal, with  production  of  a  much  higher  temperature  and 
ultimate  destruction  of  the  whole  of  the  charcoal  if  the 
volume  of  oxygen  is  large  enough.  This  union  of  the  gas 
with  the  charcoal  is  attended  by  the  escape  of  heat,  and 
evidently  intense  heat,  as  indicated  by  the  sparks.  When 
carbon  combines  with  pure  oxygen  gas,  as  in  this  case,  the 
product  formed  is  always  carbon  dioxide  and  this  substance 
may  be  always  recognized  by  the  precipitate  it  yields  with 
clear  lime-water.  It  will  be  shown  later  that  carbon  diox- 
ide is  formed  by  many  other  reactions.  , 

Other  bottles  of  the  oxygen  remain  with  which  the  fol- 
lowing experiments  may  be  made : 

Ex.  43.  Melt  a  small  amount  of  sulphur  in  a  deflagrating  spoon  (a 
small  brass  spoon  with  a  long  wire  handle  bent  to  make  a  right  angle 
with  the  rim  of  the  spoon),  heat  it  until  it  begins  to  burn  with  a  pale 
blue  flame  and  then  plunge  it  into  a  bottle  of  the  gas,  The  sulphur 
burns  here  with  a  much  brighter  flame  than  in  the  air. 

In  this  combustion  a  gas  is  produced  which  has  a  very 
strong  and  characteristic  odor.  It  is  an  oxide  of  sulphur, 
a  combination  of  oxygen  and  sulphur,  and  is  called,  prop- 
erly, sulphurous  oxide.  A  similar  but  weaker  odor  is 
noticed  when  some  kinds  of  matches  are  burned. 

The  bottle  should  be  covered  with  a  glass  plate  so  that 
the  contents  may  be  saved  for  another  experiment. 

Ex.  44.  Carefully  dry  a  "very  smah  piece  of  phosphorus  between 
folds  of  filter  paper,  place  it  in  a  deflagrating  spoon,  ignite  it  by  hold- 
ing for  a  second  in  the  lamp,  and  thrust  quickly  into  a  bottle  of  oxygen. 
Immediate  combustion  takes  place,  an  intense  white  light  is  produced, 
and  the  bottle  soon  fills  with  fumes  of  a  substance  known  as  phosphoric 
oxide.  After  a  time  these  fumes  settle  to  the  bottom  and  sides  of  the 


36  GENERAL  CHEMISTRY 

bottle  and  mix  with  the  moisture  present.     Cover  the  bottle  with  a  glass 
plate  and  keep  the  contents  for  a  second  test. 

This  experiment  gives  a  very  good  illustration  of  the 
strong  affinity ,  or  liking,  of  the  oxygen  for  phosphorus.  The 
phenomenon  is  very  similar  to  that  presented  in  the  burn- 
ing of  charcoal,  but  the  phosphorus-oxygen  reaction  is  dis- 
tinguished by  its  greater  intensity. 

The  direction  was  given  to  heat  the  phosphorus  to  the 
burning  point  before  putting  it  in  the  gas.  This  was  done 
to  save  time  rather  than  because  of  its  necessity.  If  left 
to  itself  in  dry  condition  in  the  air,  spontaneous  combustion 
takes  place  after  a  time.  Because  of  this  fact  phosphorus 
must  always  be  kept  under  water. 

In  the  experiments  just  given,  substances  which  are 
readily  combustible  have  been  burned  in  the  oxygen  gas. 
Other  bodies  burn  with  greater  difficulty,  but  in  a  manner 
none  the  less  characteristic.  This  is  true  of  many  of  the 
common  metals,  and  as  an  illustration  we  will  take  iron. 

Ex.  45.  Cut  some  fine  soft  iron  wire  into  lengths  of  about  10  cen- 
timeters afid  wrap  a  dozen  or  more  of  these  into  a  bundle;  at  one  end 
of  this  bundle  spread  the  individual  wires  a  little,  and  around  the  other 
end  twist  a  piece  of  stronger  wire  to  serve  as  a  handle.  Heat  the  loose 
end  and  dip  it  into  some  sulphur.  Ignite  this,  and  as  soon  as  it  burns 
well  remove  the  cover  from  a  jar  of  oxygen  and  hold  the  wire  down  into 
it.  First  the  sulphur  burns  with  greater  brilliancy  and  finally  the  iron 
becomes  hot  enough  to  ignite  too,  burning  in  a  manner  reminding  one 
of  the  charcoal  combustion.  The  sulphur  is  used  here  merely  to  bring 
the  iron  up  to  the  burning  temperature.  With  coarse  wire  this  experi- 
ment cannot  be  easily  performed.  In  this  combustion  the  product,  like 
the  others,  is  termed  an  oxide. 

After  this  reaction  we  have  two  substances  in  the  bottle, 
but  one  of  them  is  the  sulphurous  oxide  produced  by  the 
burning  of  the  sulphur.  The  iron  oxide  is  left  as  a  black 
scale  or  slag,  which  will  not  dissolve  in  water.  Iron  forms 
several  other  combinations  with  oxygen  which  will  be 
referred  to  later. 

It  remains  now  to  make  some  tests  with  the  products 
in  the  bottles  after  the  burning  of  the  carbon,  sulphur  and 
phosphorus.  The  oxides  of  carbon  and  sulphur  are 
invisible  gases;  that  of  phosphorus  appears  for  a  time  as  a 
white  cloud  which  subsides  as  mentioned. 


GENERAL  CHEMISTRY.  87 

Ex.  46.  To  the  contents  of  each  bottle  add  a  little  water,  replace 
the  cover  and  shake  thoroughly.  Whatever  is  present  evidently  com- 
bines with  the  water  or  goes  into  solution.  In  the  water  in  each  bottle 
dip  a  piece  of  blue  litmus  paper.  The  color  changes  to  red.  By  means 
of  a  clean  glass  rod  take  up  a  drop  of  liquid  from  each  bottle  and  touch 
it  to  the  tongue.  The  drop  from  the  carbon  bottle  imparts  little  or  no 
taste,  while  those  from  the  others  are  sour.  We  have  here  a  common 
property  of  the  bodies  usually  termed  acids,  as  most  of  these  have  a  sour 
taste  and  change  blue  litmus  paper  to  red.  The  acid  formed  by  the 
addition  of  water  to  the  carbon  bottle  is  evidently  very  weak. 

The  simple  experiments  just  detailed  are  exceedingly 
instructive.  It  can  easily  be  shown  that  oxygen  gas  itself 
is  free  from  sour  taste  and  that  it  does  not  change  the  color 
of  blue  litmus  paper.  The  addition  of  water  does  not 
alter  this.  The  acid  properties  must  therefore  come  by 
the  union  of  the  oxygen  with  the  carbon,  the  sulphur  and 
the  phosphorus.  Further,  it  can  be  shown  that  these  ox- 
ides in  the  //restate  are  not  acid  bodies;  they  become  such 
only  after  the  addition  of  water.  It  would  seem  therefore 
that  both  oxygen  and  moisture  are  concerned  in  the  pro- 
duction of  acids,  and  we  shall  find  later  that  in  very  many 
cases  this  is  true.  The  name  oxygen  signifies  acid  producer, 
it  being  at  one  time  supposed  that  all  acids  contained 
oxygen. 

The  phenomena  of  combustion  naturally  attracted  the 
attention  of  the  ancient  philosophers,  but  the  explanations 
given  of  observed  facts  were  very  faulty.  As  intimated 
above,  Lavoisier  was  the  first  to  give  a  correct  explanation 
of  what  appears  to  us  now  as  very  simple.  His  immediate 
predecessors  in  attempting  to  explain  combustion  and  oxi- 
dation were  led  to  the  doctrine  that  metals,  in  burning,  lost 
or  gave  up  something,  and  this  thing  lost  was  called  by 
them  phlogiston.  The  fact  that  these  metals  grew  heavier 
instead  of  lighter  when  burned  was  overlooked  or  consid- 
ered as  of  very  little  moment.  The  value  of  the  balance  in 
chemical  investigation  was  yet  to  be  shown.  A  few  scien- 
tific men,  recognizing,  however,  the  importance  of  the 
change  in  weight,  were  forced  to  the  absurd  conclusion  that 
phlogiston  was  endowed  with  a  property  of  levity  or  nega- 
tive weight,  so  that  when  added  to  a  body  it  made  it 
weigh  less  instead  of  more.  According  to  the  phlogiston 


38  GENERAL  CHEMISTRY. 

view  metals  were  mixtures  of  a  calx  or  base  with  phlogis- 
ton, and  application  of  high  heat  in  the  air  liberated  the 
light  phlogiston,  leaving  the  heavy  calx. 

Lavoisier  investigated  all  these  points  thoroughly  and 
established  the  fact  that  when  lead,  iron,  tin,  mercury  and 
other  common  metals  are  heated  in  the  air  they  increa'se  in 
weight  by  the  absorption  of  oxygen  ;  and  he  was  able  to 
show  in  several  cases  that  the  gain  in  weight  of  the  metal 
was  equal  to  the  loss  in  weight  of  the  volume  of  air  taken. 

He  found,  also,  as  shown  above,  that  sulphur,  phos- 
phorus and  several  other  bodies  burn  in  air  and  in  oxygen 
producing  gases,  which  in  turn  combine  with  water,  form- 
ing acids.  From  his  experiments  he  was  led  to  believe 
that  all  acids  must  contain  this  gas,  and  hence  he  gave  to 
it,  in  1778,  the  name  oxygen,  or  acid  producer.  Before 
this  date  it  had  been  known  by  several  fanciful  names,  as 
vital  air,  dephlogisticated  air,  and  others. 

The  term  vital  air  indicates  the  importance  of  oxygen 
in  the  process  of  life.  In  an  atmosphere  devoid  of  oxygen 
animals  die  almost  immediately,  and  experiment  shows 
that  a  certain  proportion  of  oxygen  must  always  be  present 
to  preserve  life.  It  has  been  mentioned  that  the  atmos- 
phere in  which  we  live  contains  about  20  per  cent  of 
oxygen.  With  a  proportion  notably  below  this,  respira- 
tion in  the  higher  animals  becomes  difficult;  if  it  falls  to  a 
certain  point,  death  must  follow.  More  will  be  said  on 
this  question  later. 

The  oxygen  taken  into  the  lungs  from  the  air  serves  for 
the  combustion  of  carbonaceous  substances  in  the  tissues. 
Here,  as  in  our  simple  experiments,  carbon  dioxide  is 
formed,  and  this  is  thrown  off  from  the  lungs  in  the  expired 
air.  As  the  carbonaceous  matters  burned  are  very  com- 
plex in  structure,  several  other  products  are  produced,  some 
of  which  are  excreted  by  the  lungs,  while  others  are  thrown 
off  by  the  kidneys  or  through  the  skin.  In  this  combustion 
within  the  body,  as  in  all  others,  heat  is  produced  and  the 
normal  high  temperature  of  animals  is  thus  seen  to  depend 
on  chemical  combination  or  oxidation. 

Conditions  of  Oxidation. — Some  substances  combine 


GENERAL  CHEMISTRY.  39 

with  oxygen  at  the  ordinary  temperature,  and  if  left 
exposed  to  the  air  burn  or  become  quickly  corroded.  The 
presence  of  moisture  in  many  cases  greatly  assists  this 
spontaneous  combination  with  oxygen,  as  is  well  illustrated 
by  the  rusting  of  iron  in  moist  air.  For  other  substances 
a  certain  elevated  temperature  must  be  reached  before 
combination  with  oxygen  will  follow.  Our  experiments 
have  shown  that  sulphur  and  charcoal  burn  readily  in 
oxygen,  but  only  after  previous  heating  to  what  is  called 
the  kindling  temperature.  It  is  a  well-known  fact  that  at 
the  ordinary  temperature  these  two  bodies  are  perfectly 
stable,  showing  no  tendency  to  oxidize. 

When  combustible  bodies  are  heated  to  the  kindling 
temperature  in  air  or  oxygen  and  burn,  a  certain  definite 
amount  of  heat  is  always  liberated.  A  gram  of  pure  car- 
bon burning  in  a  sufficient  supply  of  oxygen  liberates  a 
constant  number  of  units  of  heat,  which  term  will  be  ex- 
plained later.  The  amount  of  heat  liberated  by  the  com- 
bustion of  a  gram  of  sulphur  is  much  less  than  that  formed 
from  a  gram  of  carbon,  but  is  still  a  constant.  The  follow- 
ing short  table  gives  in  round  numbers  the  units  of  heat 
liberated  by  combustion  of  1  gram  of  the  substances  named, 
in  pure  oxygen. 

Hydrogen 34.000  units. 

Carbon 8,000 

Sulphur 2,300 

Phosphorus 5, 700 

Zinc 1,300 

Iron 1,570 

Tin 1,230 

Copper 600 

Provided  the  end  product  is  the  same  the  amount  of 
heat  liberated  is  a  constant  whether  the  oxidation  be  rapid 
or  slow,  but  the  temperature  reached  by  the  burning  body 
depends  on  the  rapidity  of  the  combustion.  In  slow  oxi- 
dation, as  in  the  rusting  of  a  piece  of  iron,  the  elevation  of 
temperature  is  not  perceptible;  but  in  rapid  combustion 
heat  enough  may  be  liberated  to  melt  the  oxide  formed. 
In  the  former  case  the  heat  liberated  is  dissipated  by  radi- 
ation or  conduction,  while  in  the  second  case  the  reaction 


40  GENERAL  CHEMISTRY. 

is  so  rapid  that  little  time  is  given  for  loss  in  this  way. 
Hence  the  elevation  of  temperature  which  follows. 

The  subject  of  the  heat  of  combustion  is  a  very  im- 
portant one  in  many  scientific  investigations,  and  also  in 
the  measurement  of  the  calorific  value  of  fuels.  A  knowl- 
edge of  the  heat  of  combustion  of  a  number  of  simple  sub- 
stances makes  it  possible  to  calculate  the  amounts  of  heat 
which  will  be  liberated  in  the  combustion  of  given  weights 
of  fuels  of  known  composition. 

One  liter  of  oxygen,  measured  at  a  temperature  of 
0°C.  and  under  a  pressure  of  760  Mm.  of  mercury  (the 
so-called  standard  temperature  and  pressure),  weighs  1.4298 
Gm.  Referred  to  dry  air  its  specific  gravity  is  1.105G.  It 
is  but  slightly  soluble  in  water,  1  volume  of  the  latter  dis- 
solving 0.041  volume  of  oxygen  at  the  standard  temperature 
and  pressure.  Oxygen  may  be  liquefied  at  a  temperature 
of  — 118°  at  a  pressure  of  50  atmospheres. 

Uses  of  Pure  Oxygen.  At  the  present  time  oxygen 
gas  as  made  by  the  chlorate  process  is  employed  for  several 
purposes.  It  is  made  for  inhalation  to  a  slight  extent,  and 
for  this  purpose  must  be  carefully  purified.  Much  larger 
quantities  are  made  for  combustion  with  hydrogen  or  illu- 
minating gas  in  the  production  of  the  "  calcium,"  'Mime" 
or  "Drummond"  light.  Oxygen  gas  is  employed  on  a 
still  larger  scale  in  several  manufacturing  operations.  Here 
it  is  made  by  other  processes  which  need  not  be  explained 
in  this  place.  It  is  now  an  article  of  commerce  and  in 
large  cities  can  be  obtained  in  any  quantity  desired  com- 
pressed in  strong  iron  cylinders. 


OZONE. 

Occurrence.  Along  with  the  oxygen  in  the  atmosphere 
several  other  so-called  oxidizing  substances  are  known  to 
exist  in  small  amount.  One  of  these  is  a  peculiar  form  of 
oxygen  itself  and  is  called  ozone.  It  must  be  remembered, 
however,  that  the  trace  of  this  substance  in  the  air  is  so 
minute  that  tests  for  its  presence  often  fail. 


GENERAL  CHEMISTRY.  41 

History.  The  peculiar  odor  noticed  in  the  neighbor- 
hood of  a  plate  electrical  machine  in  action  was  long  ago 
remarked.  In  1785  Van  Marum  called  attention  to  the 
fact  that  the  same  odor  is  developed  by  the  passage  of 
electric  sparks  through  pure  oxygen,  and  showed  that  the 
gas  so  acted  upon  has  the  power  of  immediately  tarnishing 
a  clean  surface  of  mercury.  He  found  also  that  a  decrease 
in  the  volume  of  oxygen  taken  follows.  In  1801  Cruik- 
shank  observed  that  in  the  electrolysis  of  water  acidulated 
with  sulphuric  acid  a  peculiar  odor  is  developed  at  the 
positive  pole.  Of  the  nature  of  the  substance  having  this 
marked  odor  neither  he  nor  Van  Marum  had  any  knowl- 
edge. It  was  reserved  for  the  German  chemist,  Schoen- 
bein,  to  publish  the  first  definite  details  touching  the 
production  and  properties  of  this  new  body,  which  he  did 
in  1839  and  1840.  Schoenbein  showed  that  by  the  pas- 
sage of  electricity  through  air  or  oxygen,  in  the  electrolysis 
of  water  acidulated  with  certain  acids  and  in  the  oxidation 
of  moist  phosphorus  the  same  substanceis  formed,  to  which 
he  gave  the  name  ozone.  The  true  composition  of  the  gas 
was  not  recognized  immediately.  A  number  of  important 
investigations,  extending  through  fifteen  years,  were 
required  to  fully  settle  it  to  the  satisfaction  of  all. 

Preparation.  By  special  methods  we  can  produce 
ozone  in  the  laboratory,  although  not  easily  in  large 
amount.  It  will  suffice  here  to  illustrate  this  fact  by  a 
very  simple  experiment. 

Ex.  47.  Scrape  the  surface  of  a  stick  of  phosphorus,  about  5  Cm. 
long,  under  water,  so  as  to  expose  the  pure  substance,  free  from  coating 
or  incrustation.  Next  pour  some  lukewarm  water  into  a  tall  beaker  or 
wide  mouth  bottle  to  a  depth  slightly  less  than  the  thickness  of  the 
scraped  phosphorus.  Then  transfer  the  latter  to  the  beaker  or  bottle 
and  see  that  the  surface  is  exposed  to  a  slight  extent  above  the  water  to 
the  oxidizing  action  of  the  air.  Cover  the  vessel  with  a  piece  of  glass 
and  allow  it  to  stand  some  time;  five  minutes  is  usually  sufficient. 
Meanwhile  prepare  a  test  for  the  ozone,  and  this  can  be  done  for  the 
present  instance  in  the  following  manner:  Dissolve  a  small  crystal  of 
pure  potassium  iodide  in  distilled  water  and  into  this  solution  stir  a  little 
starch.  Rub  the  starch  with  a  glass  rod  to  break  up  any  lumps  formed 
and  then  gradually  heat  to  boiling  This  makes  a  paste  in  which  the 


42  GENERAL  CHEMISTRY. 

potassium  iodide  exists  dissolved.  Into  this  paste  dip  some  small  strips 
of  filter  paper  and  then  suspend  these  in  the  vessel  in  which  the  phos- 
phorus was  left  in  contact  with  the  moist  air.  If  the  experiment  has 
been  properly  performed  the  paper  will  turn  blue,  indicating  ozone 
formed. 

The  following  explanation  must  be  given  of  the  above 
experiment.  Potassium  iodide  is  a  substance  which  is  not 
decomposed  either  in  the  dry  state  or  in  solution  by  ordi- 
nary oxygen,  but  it  is  decomposed  by  other  substances, 
among  which  is  ozone,  with  liberation  of  one  of  its  con- 
stituents— iodine.  This  iodine  forms  a  deep  blue  color 
with  starch  paste.  If,  therefore,  we  have  a  mixture  of 
potassium  iodide  and  starch  paste  and  if  from  any  cause 
this  becomes  blue,  we  know  that  some  strong  decomposing 
agent  has  acted  on  the  compound,  the  potassium  iodide, 
setting  free  its  iodine.  In  the  present  case  the  paper 
turns  blue  because  the  air  in  the  bottle  contains  a  small 
amount  of  ozone  formed  by  the  action  of  the  phosphorus 
on  the  moist  air.  The  ordinary  oxygen  has  not  the  power 
of  producing  this  blue  color. 

This  is  an  illustration  of  one  of  the  many  decomposi- 
tions which  the  ozone  gas  can  effect.  It  is  especially 
active  in  breaking  up  organic  matters,  oxidizing  or  burn- 
ing them  in  a  certain  sense.  Ozone  in  the  atmosphere  is 
supposed  to  have  the  action  of  a  purifying  agent  in  de- 
stroying decaying  animal  and  vegetable  matters. 

What  is  commonly  called  ozone  test-paper  is  prepared 
by  covering  good  book  paper  with  a  paste  of 

Water 1,000  parts 

Starch 50 

Potassium  iodide 5       " 

The  potassium  iodide  is  dissolved  in  a  small  amount  of 
the  water  in  a  porcelain  dish.  Into  this  solution  the  starch 
is  stirred  and  rubbed  with  a  pestle  until  it  forms  with  the 
water  a  uniform  creamy  liquid.  The  remainder  of  the 
water  is  then  added,  and  the  whole  is  heated,  with  con- 
stant stirring,  on  a  water-bath  until  the  starch  is  converted 
into  a  smooth  paste  free  from  lumps.  This  hot  paste  is 
spread  over  white  paper  by  means  of  a  soft  flat  brush,  and 


GENERAL  CHEMISTRY.  43 

then  the  paper  is  hung  up  to  dry  in  an  atmosphere  free 
from  oxidizing  gases.  The  dried  paper,  cut  into  small 
pieces,  may  be  kept  almost  indefinitely  in  stoppered  bot- 
tles. When  used  as  a  test  for  ozone  in  the  air,  a  small 
piece  is  moistened  and  hung  up  in  the  atmosphere  in  ques- 
tion. It  is  used  as  a  test  for  other  substances,  as  will  be 
shown  later. 

Larger  quantities  of  ozone  may  be  made  in  so-called 
ozone  generators.  These  are  forms  of  apparatus  in  which 
a  current  of  oxygen  may  be  passed  between  metallic  plates 
connected  with  the  terminals  of  an  induction  coil.  The 
silent  discharge  across  the  intervening  space  converts  a 
part  of  the  gas  into  ozone.  But  the  reaction  is  always 
far  from  complete,  unless  the  product  is  absorbed  by  po- 
tassium iodide  solution,  or  something  else,  as  fast  as 
formed. 

Ozone  has  been  shown  to  be  formed  by  the  conden- 
sation of  ordinary  oxygen  in  a  peculiar  manner,  which  will 
be  referred  to  later.  In  this  condensation  3  volumes  of 
oxygen  yield  2  of  ozone.  At  a  temperature  of  300°  C.  this 
condensed  product  is  completely  decomposed,  common 
oxygen  resulting. 

The  oxidizing  action  of  ozone  is  powerful,  many 
organic  substances  being  quickly  destroyed  by  it.  As  a 
bleaching  agent  it  is  many  times  as  strong  as  chlorine.  In 
the  older  literature  (since  1850)  it  was  considered  as  the 
most  powerful  natural  purifying  agent  in  the  atmosphere, 
but  it  is  now  generally  admitted  that  most  of  the  effects 
ascribed  to  ozone  in  the  air  are  due  to  a  related  body,  the 
peroxide  of  hydrogen,  which  will  be  described  later,  or 
to  nitrous  acid,  which  is  present  in  small  traces. 


HYDROGEN. 

This  is  a  gaseous  element  of  the  highest  importance 
from  many  standpoints. 

Occurrence.     As  a  free  substance  it  is  found  in  nature 
in  traces  only,  but  is  one  of  the  common  elements  in  com- 


44  GENERAL  CHEMISTRY. 

bination,  constituting  one-ninth  by  weight  of  water  and  an 
important  fraction  of  most  animal  and  vegetable  sub- 
stances. 

History.  When  iron  and  zinc  are  dissolved  in  dilute 
acids  an  inflammable  gas  is  evolved.  This  fact  was 
observed  by  some  of  the  alchemists,  but  received  no  expla- 
nation from  them  or  their  followers.  Cavendish,  in  1766, 
published  an  investigation  of  the  subject  in  which  he  de- 
scribes the  gas  as  inflammable  air.  Later  he  considered  it 
as  identical  with  pure  phlogiston,  because  it  was  found 
capable  of  regenerating  pure  metals  from  the  calces  referred 
to  under  oxygen,  and  this  view  was  held  by  many  others. 
Following  up  his  experiments,  Cavendish  found  with  con- 
siderable accuracy  the  amount  of  the  gas  which  may  be 
liberated  from  acids  by  given  weights  of  several  metals. 
In  1781  he  found  that  water  is  composed  of  inflammable 
air  and  dephlogisticated  air,  but  at  the  time  he  apparently 
failed  to  realize  the  importance  of  his  discovery.  In  1783 
Lavoisier  gave  the  first  clear  explanation  of  the  composi- 
tion of  water  and  proposed  the  name  hydrogen  for  the 
inflammable  air  of  Cavendish.  About  this  time  it  was 
found  that  hydrogen  may  be  obtained  by  the  action  of  cer- 
tain metals  on  strong  alkali  solutions  as  well  as  on  acids. 

Preparation.  We  are  able  to  separate  hydrogen  from 
its  compounds  by  many  simple  reactions.  In  illustration 
of  these  we  will  consider  first  the  decomposition  of  water. 
It  has  been  stated  before  that  water  is  a  compound  of 
hydrogen  and  oxygen,  and  in  separating  them  we  must 
overcome  the  strong  affinity  which  holds  them  together. 
This  may  be  conveniently  done  bypassing  a  strong  cur- 
rent of  electricity  through  the  water  slightly  acidulated 
with  sulphuric  acid.  Certain  metals  brought  in  contact 
with  water  are  also  able  to  separate  the  hydrogen  through 
the  attraction  they  have  for  the  oxygen.  The  following 
experiment  will  show  how  this  separation  may  be  effected 
at  the  ordinary  temperature: 

Ex.  48.  Drop  a  small  piece  of  sodium,  not  larger  than  a  pea,  on 
the  surface  of  water  contained  in  an  earthenware  bowl.  As  the  metal  is 


GENERAL  CHEMISTRY.  45 

lighter  than  water  it  floats,  but  as  it  does  so  a  decomposition  goes  on  in- 
dicated by  the  escape  of  a  gas  with  a  sound  reminding  one  of  the  escape 
of  steam.  On  striking  the  water  the  sodium  melts  and  assumes  the 
globular  form.  It  soon  becomes  evident  that  in  the  reaction  between 
the  two  the  sodium  is  worn  away,  as  the  globule  grows  small  and  finally 
disappears.  While  the  gas  is  escaping  bring  a  small  flame  in  contact 
with  it  and  observe  that  it  ignites  and  burns  readily  with  a  yellow  color. 
Very  often  the  gas  ignites  spontaneously.  In  performing  this  experi- 
ment it  sometimes  happens  that  the  sodium  globule  flies  into  small  bits 
which  are  scattered  in  all  directions.  The  face  should  not  be  held  over 
the  bowl,  therefore,  when  making  the  test. 

Sodium  is  a  light,  silver  white  metal  which  is  kept 
under  rock  oil.  When  a  piece  is  taken  out  for  this  experi- 
ment it  is  cut  to  the  proper  size  and  wiped  free  from  the 
oil  by  means  of  filter  paper.  The  liberated  gas  may  be 
collected  in  a  bottle  or  test-tube  and  examined.  To  do 
this,  fill  the  bottle  or  tube  with  water  and  invert  it  in 
the  bowl  in  the  usual  manner.  Then  wrap  a  bit  of  clean 
dry  sodium  in  filter  paper  or  wire  gauze,  and  by  means  of 
tongs  bring  the  pellet  so  enclosed  under  the  mouth  of  the 
bottle  or  tube,  held  for  the  purpose  just  below  the  surface 
of  the  water  in  the  bowl.  The  decomposition  takes  place 
as  before,  but  the  gas  ascends  into  the  receptacle,  and  may 
be  tested  as  explained  later.  The  beginner  is  not  advised 
to  make  this  experiment.  The  yellow  color  of  the  flame  is 
not  characteristic  of  hydrogen,  but  of  the  vapor  of  sodium 
burned  with  it. 

Ex.  49.  Repeat  the  last  experiment,  using  a  small  piece  of  potas- 
sium instead  of  sodium.  Observe  the  same  precautions.  The  flame  is 
now  violet. 

Ill  both  of  these  experiments  hydrogen  gas  is  set  free 
and  a  certain  amount  of  water  has  been  converted  into 
something  else.  The  character  of  this  is  disclosed  by  two 
simple  tests.  First,  by  means  of  a  glass  rod  touch  a  little 
of  the  water  to  the  tongue;  a  sharp  caustic  taste  will  be 
noticed.  Then  dip  a  piece  of  red  litmus  paper  in  it  and 
notice  the  change  of  color  to  blue.  This  is  evidence  of  the 
presence  of  alkali.  The  same  evidence  is  given  by  the 
deep  red  color  produced  when  a  few  drops  of  alcoholic 
phenol-phthalein  solution  are  added  to  the  water.  It 
appears,  therefore,  that  in  this  experiment  with  the  water 


46 


GENERAL  CHEMISTRY. 


hydrogen  gas  is    liberated,  while  an    alkali   substance  is 
formed. 

Sodium  and  potassium  are  not  the  only  metals  which 
decompose  water  in  the  cold,  while  at  a  high  temperature 
the  reaction  is  possible  with  many  others,  as  is  readily 
shown.  The  usual  method  of  procuring  hydrogen  for  ex- 
periment depends  on  the  decomposition  of  some  acid  by 
means  of  a  metal.  In  Ex.  29  it  was  shown  that  when 
dilute  sulphuric  acid  is  poured  over  zinc  in  a  beaker  the 


FIG.  G. 

metal  gradually  dissolves  with  evolution  of  gas.  This  gas 
is  hydrogen  and  the  experiment  may  now  be  arranged  to 
collect  and  test  it. 

Ex.  50.  Arrange  a  gas  generating  bottle  as  shown  in  the  cut  above. 
The  bottle  should  hold  about  250  Cc.  and  be  closed  with  a  doubly  per- 
forated stopper.  Through  one  perforation  the  stem  of  a  funnel  tube 
passes  while  the  gas  is  led  out  through  a  "delivery  tube"  from  the 
other.  Put  some  granulated  zinc  in  the  bottle,  insert  the  stopper  with 
the  funnel  tube  and  delivery  tube,  add  some  water  and  then  some  dilute 
sulphuric  acid  so  as  to  about  one-third  fill  the  bottle  and  cover  the 
lower  end  of  the  funnel  tube.  An  evolution  of  gas  soon  begins,  and  if 
the  end  of  the  delivery  tube  is  brought  under  water  the  gas  passes 
through  and  may  be  collected  as  in  the  case  of  oxygen.  Therefore 
arrange  several  bottles  for  the  collection  of  gas  as  there  described,  and 
after  it  has  bubbled  through  the  water  a  few  minutes,  allow  them  to 


GENERAL  CHEMISTRY.  4? 

fill  with  the  gas.  When  filled  remove  them  by  aid  of  a  glass  plate  but 
keep  the  mouths  of  the  bottles  now  down,  because  the  gas  is  much 
lighter  than  air.  When  several  bottles  are  filled  their  contents  may  be 
tested. 

Ex.  51.  Lift  one  of  the  bottles  from  the  table,  mouth  still 
down,  and  thrust  up  into  it  a  lighted  taper  or  splinter  of  burning  wood. 
As  the  light  goes  into  the  bottle  it  is  extinguished  but  a  flame  appears  at 
the  mouth  of  the  bottle,  from  the  ignited  hydrogen.  The  gas,  there- 
fore, burns,  but  the  combustion  on  the  end  of  the  taper  or  splinter  is 
checked,  because,  as  shown  before,  oxygen  is  necessary  for  this  and  the 
bottle  contains  hydrogen.  The  hydrogen  itself,  at  the  mouth  of  the  bot- 
tle, enters  into  combination  with  the  oxygen  of  the  air. 

Ex.  52.  Invert  one  of  the  filled  bottles,  holding  the  mouth  now  up, 
remove  the  glass  plate  and  bring  a  flame  to  the  gas.  It  will  ignite  with 
a  slight  report  and  burn  in  the  bottle,  as  the  heavier  oxygen  of  the  air 
tends  now  to  settle  and  mix  with  the  hydrogen.  In  nearly  all  cases  the 
flame  shows  some  color.  This  is  not  characteristic  of  the  hydrogen  but  of 
various  impurities  with  it  or  on  the  surface  of  the  glass.  The  flame 
from  pure  hydrogen  has  a  slight  blue  tinge  only. 

Ex.  53.  The  relative  lightness  of  hydrogen  may  be  readily  shown 
by  pouring  it  upward  into  a  bottle  filled  with  air.  Use  one  of  the  bot- 
tles of  hydrogen  still  standing  on  a  glass  plate.  Lift  it  with  one  hand, 
and  turn  it  so  that  its  mouth  is  brought  under  and  near  that  of  the  air- 
filled  bottle  of  the  same  size.  After  a  minute  or  two  hold  the  upper 
bottle  to  a  flame,  when  a  sharp  report  shows  the  presence  of  hydrogen. 
The  lower  bottle  still  contains  some  of  the  gas,  which  can  be  shown  in 
the  same  manner. 

It  will  be  seen  later  that  certain  gases  heavier  than  air 
can  be  poured  downward  from  one  vessel  into  another,  as 
in  the  case  of  water.  This  is  true  of  carbonic  acid  gas  and 
chlorine,  for  instances.  In  the  above  experiment  the  light 
hydrogen  ascends,  and  a  certain  amount  of  the  air  in  the 
upper  bottle  is  forced  out  by  it.  The  displacement  is  not 
perfect,  however,  as  was  shown  by  the  manner  of  the  ex- 
plosion of  the  gaseous  mixtures  in  the  two  bottles  above. 

Hydrogen  from  Other  Sources.  While  hydrogen 
is  commonly  prepared  as  just  shown,  it  may  be  made  by 
the  action  of  certain  metals  on  alkali  solutions.  Aluminum 
wire,  for  instance,  decomposes  a  solution  of  potassium  or 
sodium  hydroxide  very  readily,  especially  when  aided  by 
heat.  The  reaction  may  be  carried  out  in  a  test-tube,  and 


48 


GENERAL  CHEMISTRY. 


the  character  of  the  escaping  gas  determined.  The  same 
alkali  solutions  are  decomposed  by  other  metals  also. 

These  reactions  have  practical  value,  as  in  certain 
investigations  it  is  desirable  to  liberate  hydrogen  without 
the  use  of  an  acid,  and  the  alkali  methods  may  then  often 
be  applied. 

Hydrogen  is  also  easily  liberated  from  water  by  the 
passage  of  the  electric  current,  as  intimated.  This  will  be 
illustrated  later  by  an  experiment  to  determine  the  com- 
position of  water. 

It  has  been  shown  that  sodium  and  potassium  decom- 
pose cold  water  with  liberation  of  hydrogen.  At  a  higher 
temperature  the  same  decomposition  may  be  effected  by 
other  metals.  When  steam  is  passed  through  an  iron  or 


FIG.  7. 


porcelain  tube  containing  iron  turnings,  and  heated  to  a 
very  high  temperature  in  a  gas  furnace,  it  is  decomposed, 
the  hydrogen  being  set  free,  while  the  oxygen  remains  in 
combination  with  the  iron,  forming  an  oxide  of  iron.  This 
experiment  is  easily  carried  out  in  the  apparatus  illustrated 
by  the  above  figure.  Water  is  boiled  in  a  flask  to  the  left. 
The  steam  generated  passes  into  the  tube  resting  over  a 
number  of  burners  in  the  furnace,  while  the  liberated 
hydrogen  is  collected  in  a  jar  beyond.  By  filling  the  tube 
with  charcoal  instead  of  with  iron  turnings,  a  somewhat 
analogous  decomposition  takes  place.  We  obtain  now 
hydrogen  mixed  with  oxides  of  carbon,  as  the  carbon  com- 
bines with  the  oxygen  of  the  water  to  form  these  bodies, 
which  are  gases. 


GENERAL  CHEMISTRY,  49 

The  last  decomposition  is  a  very  important  one,  as  it 
is  the  basis  of  the  process  commonly  followed  in  the  man- 
ufacture of  water  gas,  generally  used  at  the  present  time. 
It  will  be  fully  described  later. 


Diffusion  of  Hydrogen.  Because  of  its  extreme  light- 
ness this  gas  is  very  suitable  for  showing  an  interesting 
property  of  all  gases,  viz. :  that  of  diffusion.  Two  gases 
separated  from  each  other  by  a  porous  partition — a  thin 
plate  of  plaster  of  Paris,  for  instance — will  in  time  mix 
with  each  other,  as  both  pass  through  the  porous  sub- 
stance. The  rates  of  diffusion  or  passage  of  the  gases 
bear  a  close  relation  to  their  specific  gravities  or  densities. 
It  has  been  found  that  the  velocity  of  diffusion  is  inversely 
proportional  to  the  square  root  of  the  density  of  a  gas. 
From  this  it  would  follow  that  hydrogen  must  diffuse  4 
times  as  fast  as  oxygen  and  3.8  times  as  fast  as  air,  as  the 
densities  of  the  gases  stand  to  each  other  in  the  relation, 
1  :  16  :  14.45.  A  simple  proof  that  hydrogen  moves  much 
more  rapidly  than  air  is  given  in  the  following  experiment: 

Ex.  54.  From  a  piece  of  glass  tubing  having  an  internal  diameter 
of  a  centimeter  or  more,  cut  off  a  length  of  about  twenty  centimeters. 
Dip  one  end  of  this  into  some  soft  plaster  of  Paris,  so  as  to  take  up  a 
plug  about  one  centimeter  in  thickness.  Set  the  tube  aside  for  this  to 
harden,  which  will  require  some  hours  Then  fill  it  with  hydrogen  gas 
by  displacement  of  the  air,  and  immediately  stand  it  in  upright  position 
in  a  beaker  of  water  in  such  a  position  that  the  open  end  is  covered  to  a 
depth  of  several  centimeters.  It  will  be  observed  that  the  water  ascends 
in  the  tube  and  finally  reaches  a  position  much  above  the  level  of  the 
water  in  the  beaker.  It  then  recedes  slowly  and  in  time  the  level  cor- 
responds to  that  outside. 

In  the  first  stage  of  this  experiment  the  water  ascends 
because  the  hydrogen  passes  out  through  the  porous  plug 
much  faster  than  the  air  can  enter.  The  maximum  posi- 
tion of  the  water  is  reached  when  the  rates  of  motion  in 
opposite  directions  are  equal,  after  which  the  column  of 
water  falls  because  the  volume  of  gas  entering  the  tube  is 
now  greater  than  that  leaving  it.  Many  porous  stones  may 
be  used  to  exhibit  this  phenomenon. 


50  GENERAL  CHEMISTRY. 

Reducing  Power  of  Hydrogen.  By  this  we  under- 
stand the  property  which  hydrogen  possesses  of  abstract- 
ing oxygen  from  certain  compounds,  forming  with  it  water. 
The  term  is  used  also  in  a  broader  sense  but  in  this  place 
the  limited  usage  only  will  be  considered. 

It  will  be  shown  later  by  experiment  than  when  hydro- 
gen gas  is  passed  over  the  oxides  of  copper  or  iron, heated 
to  a  high  temperature,  the  oxygen  is  taken  and  the  metal 
left  in  the  free  state.  Other  metallic  oxides  may  be 
reduced  in  a  similar  manner. 

This  reaction  is  one  of  the  highest  importance  and  is 
frequently  employed  in  the  laboratory  for  several  purposes. 
An  illustration  will  be  given  in  the  next  section. 

An  animal  placed  in  an  atmosphere  of  hydrogen  would 
soon  die,  but  this  would  follow  from  asphyxiation  rather 
than  from  any  poisonous  property  of  the  gas.  In  an  atmos- 
phere of  4  parts  of  hydrogen  and  1  part  of  oxygen  animals 
live  apparently  as  well  as  in  ordinary  air. 

Hydrogen  is  but  slightly  soluble  in  water,  1  volume  of 
the  latter  dissolving  of  the  gas  0.0193  volume  at  a  tem- 
perature of  0°C.,  and  under  a  pressure  of  760  Mm.  of 
mercury.  Several  metals  possess  the  power  of  absorbing 
hydrogen  in  considerable  quantity.  In  the  case  of  the 
metal  palladium  this  power  is  very  marked.  The  cold 
metal  absorbs  about  375  times  its  volume  of  hydrogen, 
while  at  a  red  heat  nearly  1,000  volumes  are  absorbed. 

Under  standard  conditions  1  liter  of  pure  hydrogen 
weighs  0.0900  Gm.  It  has  been  condensed  to  a  colorless 
liquid  at  a  temperature  below  — 200°C.  with  a  pressure  of 
40  atmospheres. 

WATER. 

Occurrence.  The  student  is  familiar  with  the  natural 
occurrence  of  water  in  the  seas,  lakes,  rivers,  etc.  The 
purification  of  water  by  distillation  has  been  referred  to 
already,  and  further  details  of  practical  processes  will  be 
given  later. 

History.     In  the  preceding  section  it  was  explained 


GENERAL  CHEMISTRY  51 

that  the  exact  composition  of  water  was  first  suggested  by 
the  experiments  of  Cavendish,  while  Lavoisier's  work 
proved  the  fact  conclusively.  Careful  investigations  un- 
dertaken by  Gay  Lussac  and  Humboldt  were  published  in 
1805,  and  these  confirmed  the  work  of  Lavoisier  and 
showed  that  exactly  two  volumes  of  hydrogen  combine 
with  1  volume  of  oxygen  to  form  water. 

Some  experiments  will  now  be  given  to  illustrate 
methods  of  finding  this  ratio. 

Composition  of  Water.  The  presence  of  hydrogen 
in  water  was  suggested  by  the  experiment  in  which  metal- 
lic sodium  was  used  to  decompose  water.  Other  metals, 
as  intimated,  behave  in  a  similar  manner.  At  a  high  tem- 
perature, water  is  readily  decomposed  by  iron  turnings 
with  liberation  of  hydrogen,  which  has  been  referred  to 
already.  A  still  more  convenient  method  of  decomposition 
is  by  means  of  the  electric  current,  which  will  be  experi- 
mentally shown. 

Water  in  absolutely  pure  condition  is  not  a  conductor 
of  electricity  and  therefore  in  this  form  is  not  decomposed 
by  it.  But  by  the  addition  of  a  little  acid  to  water  it  be- 
comes a  moderately  good  conductor,  and  the  current  which 
may  now  be  made  to  pass  through  effects  decomposition. 
Acids  are  not  the  only  substances  which  when  added  to 
water  render  it  a  conductor,  but  for  our  purpose  they  are 
the  most  convenient.  If  two  plates  or  strips  of  thin  plati- 
num foil,  attached  to  the  opposite  poles  of  a  galvanic  bat- 
tery of  several  Bunsen  or  Leclanche"  cells,  be  dipped  in  a 
beaker  of  acidulated  water,  gas  bubbles  will  be  seen  to 
ascend  from  the  surface  of  each  plate.  If  these  plates  be 
supported  beneath  two  tubes  filled  with  water  the  gas  bub- 
bles will  pass  up  into  them  and  displace  the  water.  After 
a  time  the  contents  of  each  tube  may  be  tested.  If  the 
tubes  are  so  placed  that  they  collect  all  the  gas  given  off 
from  each  plate  in  a  given  time  it  will  be  noticed  that  one 
volume  is  almost  exactly  twice  the  other,  and  that  the 
tests  of  the  larger  volume  show  it  to  be  hydrogen,  while 
the  smaller  volume  gives  the  tests  characteristic  of  oxygen. 
Several  special  forms  of  apparatus  have  been  devised  for 


52 


GENERAL  CHEMISTRY 


this  experiment  of   which  the  one  now  to  be  described  is 
very  convenient. 

Ex.  55.  Arrange  the  apparatus  as  shown  in  the  next  figure  (Fig. 
8).  As  seen  it  consists  essentially  of  a  long  U  tube,  the  two  limbs  of 
which  are  closed  on  top  by  ground  glass  stopcocks.  From  the  bottom 
of  the  bend  a  tube  passes  backward  and  then  upward,  ending  finally  in 
a  wide  bulb  or  reservoir.  Inside  of  each  limb  of  the  U  tube,  just  above 
the  bend,  there  is  a  thin,  platinum  plate,  which  is  connected  with  a 
short  platinum  wire  passing  through  the  glass  and  ending  in  a  loop  just 
outside. 


FJG.  8. 


By  opening  the  stopcocks  on  the  U  tube,  the  apparatus  may  be 
filled  with  acidulated  water,  to  this  level,  by  pouring  into  the  bulb  tube. 
The  stopcocks  are  then  closed.  (Water  containing  about  5  per  cent  of 
sulphuric  acid  is  suitable  for  the  purpose.)  The  apparatus  is  now  ready 
for  the  actual  experiment  which  is  begun  by  attaching  the  copper  wires 
from  a  good  battery  to  the  platinum  loops  described  above.  Almost 
immediately  gas  bubbles  form  on  each  plate  and  escape  up  into  the 
tubes  above.  As  the  plates  are  situated  above  the  bend  of  the  U,  the 
gases  cannot  mix  on  liberation,  but  must  remain  separate.  In  a  few 
minutes  it  becomes  apparent  that  the  volume  of  the  gas  liberated 
from  the  plate  connected  with  the  negative  pole  of  the  battery,  the  pole 
connected  with  the  zinc,  is  practically  twice  the  volume  liberated  at  the 
other  pole. 


GENERAL  CHEMISTRY.  53 

Allow  the  experiment  to  continue  until  the  larger  volume  fills  the 
limb  down  to  the  platinum  plate.  It  will  be  observed  that  as  the  gases 
collect  the  liquid  is  forced  down  and  then  up  into  the  bulb  tube,  and 
further  that  the  diminution  of  the  volume  of  liquid  itself  is  not  great. 
In  fact,  not  more  than  a  small  drop  of  the  liquid  undergoes  decomposi- 
tion to  form  the  relatively  large  gas  volume. 

The  nature  of  the  gases  in  the  two  limbs  may  now  be  tested.  To  this 
end  wipe  off  the  tips  above  the  stopcocks,  and  free  them  as  far  as  pos- 
sible from  liquid  by  means  of  bibulous  paper,  then  light  a  small  taper  or 
splinter,  hold  it  over  the  limb  with  the  larger  volume  and  carefully  open 
the  stopcock.  The  pressure  of  the  liquid  in  the  bulb  tube  will  force  the 
gas  out,  and  this  burns  in  a  manner  characteristic  of  hydrogen,  on  com- 
ing in  contact  with  the  small  flame.  Over  the  other  limb  hold  a  glow- 
ing splinter.  On  opening  the  stopcock  the  gas  which  streams  out  gives 
the  behavior  characteristic  of  oxygen.  According  to  this  experiment  it 
would  appear  that  water  may  be  decomposed  into  two  volumes  of  hydro- 
gen and  one  of  oxygen. 

By  repeating  the  above  experiment  carefully  with  ac- 
curate apparatus  it  can  be  shown  that  one  volume  is  a  lit- 
tle more  than  twice  the  other.  The  oxygen  volume  is  rel- 
atively small  because  this  gas  is  more  soluble  in  the  liquid 
than  is  the  hydrogen,  and  also  because  a  little  of  it  is 
changed  into  ozone  in  the  reaction.  Both  of  these  facts 
have  been  referred  to  already.  If  precautions  are  taken  to 
avoid  the  production  of  ozone  and  if  the  solubility  of  the 
oxygen  is  diminished  by  working  at  a  high  temperature 
then  the  two  volumes  will  be  found  to  be  liberated  in  the 
exact  proportions  of  2:1. 

As  carried  out,  the  last  experimentcould  not  be  regarded 
as  conclusive,  as  not  all  of  the  liquid  was  decomposed. 
But  the  experiment  may  be  repeated  as  often  as  desired 
with  the  water  remaining  after  each  test  until  its  volume 
becomes  quite  small.  The  result  of  the  electrolysis  is  the 
same  in  all  cases. 

Further  information  concerning  the  composition  of 
water  is  given  by  the  following  experiments. 

If,  as  suggested  by  the  foregoing,  water  is  composed  of 
hydrogen  and  oxygen  gases  we  should  be  able  to  produce 
water  by  the  combination  of  these  two  substances.  The 
experiment  may  be  readily  made  and  according  to  several 
plans. 

Ex.  56.     First,  we  may  make  a  very  simple  experiment  with  the 


54 


GENERAL  CHEMISTRY. 


apparatus  illustrated  below.  To  the  common  hydrogen  generator  is 
attached  a  so-called  drying  tube,  which  contains  some  substance  to 
absorb  moisture  from  the  gas.  Calcium  chloride  is  often  used  for  the 
purpose.  The  generated  gas  after  passing  through  the  drying  tube 
reaches  the  air  by  means  of  a  narrow  tube  with  a  fine  opening.  The 
hydrogen  gas,  after  having  passed  long  enough  to  expel  the  air,  is  lighted 
at  this  opening  and  burns  with  a  small  hot  flame.  If  now  a  cold  dry 
beaker  be  held,  mouth  down,  over  this  flame  a  deposit  of  moisture 
immediately  collects  on  the  cold  surface  of  the  glass.  This  must  be 
formed  by  a  union  of  the  hydrogen  with  something  in  the  air. 


FIG.  9. 


That  this  substance  from  the  air  is  the  oxygen  may  be 
shown  by  trial  with  the  pure  gas. 

Given  volumes  of  the  two  gases  may  be  combined 
directly  by  aid  of  an  electric  spark  and  the  result  noted, 
when  it  will  be  found  that  exactly  two  volumes  of  hydrogen 
unite  with  one  volume  of  oxygen.  For  our  purpose  it  will 
prove  more  convenient  to  make  the  determination  in 
another  manner,  somewhat  less  direct,  but  equally  con- 
clusive. Instead  of  combining  the  hydrogen  with  pure 
oxygen  we  can  allow  hydrogen  gas  to  act  on  some  sub- 


GENERAL  CHEMISTRY. 


55 


stance  containing  oxygen  in  combination,  under  such  con- 
ditions that  the  substance  will  decompose,  giving  up  its 
oxygen  to  the  hydrogen.  In  such  a  case  hydrogen  acts  as 
a  reducing  agent,  as  was  explained  some  pages  back,  and 
the  oxides  of  copper  and  iron  are  good  illustrations  of  sub- 
stances which  may  be  decomposed  in  this  manner.  The 
oxide  of  copper  is  a  compound  easily  made  in  condition  of 
great  purity  so  that  there  is  no  doubt  regarding  its  exact 
composition  and  the  proportion  of  oxygen  it  contains.  The 
experiment  which  the  student  is  able  to  make  with  this 
substance  is  a  qualitative  one,  but  the  modifications  neces- 
sary to  make  it  quantitative  will  be  suggested. 


FIG.  10. 


Ex.  57.  Arrange  the  apparatus  as  shown  in  the  above  figure.  It 
consists  of  a  hydrogen  generator  and  a  tube  following  it  filled  with  cal- 
cium chloride  to  dry  the  gas.  Then  follows  a  piece  of  hard  glass  tub- 
ing, plain  or  furnished  with  a  bulb,  to  contain  the  pure  dry  copper 
oxide  to  be  decomposed.  After  the  tube  of  hard  glass  comes  a  bent,  or 
U  tube,  to  catch  the  products  of  the  reaction  to  be  described.  This  U 
tube  is  immersed  in  a  beaker  of  cold  water,  and  in  turn  is  joined  to  a 
second  small  calcium  chloride  tube.  The  connections  must  all  be  made 
with  perfect,  sound  corks,  which  fit  with  absolute  accuracy.  The  appa- 
ratus being  ready,  and  the  hard  glass  tube  charged  with  10  to  20  Gm.  of 
the  copper  oxide,  we  begin  the  experiment  by  generating  hydrogen 
in  the  usual  manner  by  the  action  of  dilute  sulphuric  acid  on  zinc.  The 


56  GENERAL  CHEMISTRY. 

gas,  given  off  in  the  moist  condition,  is  dried  as  it  passes  through  the 
calcium  chloride  tube  and  then  enters  the  combustion  tube,  from  which 
it  displaces  the  air.  After  the  gas  has  passed  some  minutes  through  the 
apparatus  the  flame  of  a  Bunsen  burner  is  applied  to  the  combustion 
tube  immediately  under  the  portion  of  the  oxide  of  copper  nearest  the 
generator.  The  heat  must  be  very  carefully  applied  at  first  to  avoid 
cracking  the  hard  tube,  which  can  be  best  prevented  by  having  the  flame 
low  and  by  moving  it  to  and  fro  along  the  tube.  After  a  few  minutes 
stronger  heat  may  be  applied .  It  will  soon  be  noticed  that  vapor  ascends 
from  the  black  mass  and  that  it  is  pushed  forward  by  the  pressure  of  the 
hydrogen  toward  the  U  tube,  and  also  that  in  a  short  time  the  heated 
oxide  above  the  flame  glows  as  if  on  fire.  We  have  here  the  stage  of 
active  reduction.  It  becomes  presently  evident  that  in  this  very  hot 
part  of  the  tube  the  oxide  of  copper  has  been  converted  into  bright 
metallic  copper,  the  color  of  which  is  very  distinct.  When  the  reaction 
here  is  complete  the  flame  is  moved  onward  toward  the  U  tube  and  the 
operation  continued  until  the  whole  of  the  black  oxide  has  undergone 
decomposition.  The  vapors  given  off  suggest  the  formation  of  water, 
but  this  can  be  further  tested  in  the  U  tube.  These  vapors  tend  to  con- 
dense in  the  cooler  part  of  the  combustion  tube  itself,  but  the  condensed 
product  can  be  dislodged  by  careful  application  of  the  lamp  heat  and 
completely  driven  over  into  the  U  tube  in  the  beaker  of  water.  A  small 
portion  of  the  vapor  passes  through  the  U  tube,  but  is  intercepted  by  the 
calcium  chloride  tube  beyond.  After  the  completion  of  the  reduction, 
that  is,  when  the  black  oxide  of  copper  has  disappeared,  leaving  only 
bright  metal,  the  heat  is  withdrawn,  leaving  the  tube  to  cool  down  with 
the  current  of  hydrogen  still  passing.  Then  it  is  taken  apart  and  the 
contents  of  the  U  tube  examined.  The  proper  tests,  which  the  student 
is  not  prepared  to  make  at  this  stage  of  his  work,  however,  show  that  the 
condensed  liquid  is  water,  and  nothing  else.  It  has  evidently  been 
formed  by  the  union  of  the  hydrogen  gas  with  the  oxygen  taken,  as 
needed,  from  the  copper  oxide. 

It  sometimes  happens  in  this  experiment  that  the  water 
collected  is  slightly  colored.  This  is  due  to  the  action  of 
heat  on  the  corks  at  each  end  of  the  combustion  tube,  and 
can  be  avoided  by  careful  manipulation.  The  oxide  of 
copper  should  be  pure  and  free  from  any  kind  of  organic 
dust,  and  should  occupy  the  central  portion  of  the  combus- 
tion tube.  This  must  be  quite  clean  and  dry  at  the  begin- 
ning of  the  experiment.  The  oxide  of  copper  may  be  held 
in  proper  place  by  means  of  a  loose  tuft  of  clean  glass  wool 
at  each  end. 

To  make  the  experiment  perfect  in  all  details  would 
necessitate  the  complication  of  the  apparatus  to  a  degree 
which  would  render  it  unfit  for  our  present  purpose.  As 


GENERAL  CHEMISTRY.  57 

constructed  and  manipulated  above  it  is  sufficient  to  show 
conclusively  that  only  hydrogen  and  oxygen  are  concerned 
in  the  formation  of  water.  By  the  introduction  of  certain 
modifications  the  experiment  may  be  made  a  quantitative 
one.  It  is  simply  necessary  to  provide  more  perfect  ap- 
paratus to  absorb  the  water  formed,  and  weigh  this  accu- 
rately before  the  experiment.  It  is  weighed  again  at  the 
end,  the  increase  in  weight  showing  exactly  the  amount  of 
water  produced.  The  oxide  of  copper  must  be  weighed 
before  the  experiment,  and  the  residue  left  after  it.  The 
loss  of  weight  here  corresponds  to  the  oxygen  given  to  the 
hydrogen.  The  experiment, therefore,  gives  us  the  weight 
of  oxygen  in  a  determined  weight  of  water;  the  difference 
between  these  two  weights  represents  the  hydrogen,  be- 
cause nothing  else  has  been  used  in  the  work. 

This  is  a  fundamental  experiment  upon  which  many 
skilled  chemists  have  spent  a  great  deal  of  time,  for  the 
purpose  of  determining  with  the  highest  possible  accuracy 
the  exact  ratio  in  which  hydrogen  and  oxygen  combine. 
This  knowledge,  as  will  appear  later,  is  of  great  scientific 
importance,  as  much  depends  on  it. 

A  very  satisfactory  method  of  showing  the  direct  union 
of  the  two  gases  in  the  proportion  mentioned,  is  by  means 
of  an  explosion  of  a  mixture  of  the  gases  in  an  apparatus 
known  as  an  eudiometer,  shown  in  the  next  figure.  This 
apparatus  consists  essentially  of  a  long,  accurately  grad- 
uated glass  tube,  the  divisions  being  usually  in  millimeters. 
Near  the  closed  end  of  the  tube  two  pieces  of  platinum 
wire  are  passed  through  the  glass  and  sealed  in  so  that 
they  nearly,  but  not  quite,  touch  inside  the  tube. 

The  eudiometer  is  filled  with  mercury,  inverted  in  a 
reservoir  of  mercury,  b,  and  clamped  in  position.  A  certain 
volume  of  pure  hydrogen  is  now  allowed  to  enter  the  tube 
and  is  accurately  measured,  the  tube  being  brought  to  a 
vertical  position  for  this  purpose,  and  for  all  subsequent 
measurements.  The  illustration  shows  the  tube  in  an 
inclined  position,  into  which  it  is  brought  for  convenience 
in  filling. 

A  volume  of  pure  oxygen  over  half  as  great  is  next 
added  and  the  new  mixed  volume  accurately  noted.  The 


58  GENERAL  CHEMISTRY. 

open  end  of  the  eudiometer  is  pressed  down  on  a  piece  of 
rubber  under  the  mercury  and  then  the  tube  is  firmly 
clamped.  A  spark  from  an  induction  coil  is  caused  to 
jump  between  the  wires  within  the  tube.  This  produces 
an  explosion  in  the  gaseous  mixture,  and  when  the  pres- 
sure on  the  rubber  plate  is  released  it  will  be  noticed  that 
mercury  rushes  up  into  the  tube.  After  a  time  the  remain- 
ing gas  volume  is  accurately  measured.  A  subsequent 
test  shows  that  it  is  oxygen.  If  this  volume  is  subtracted 
from  the  oxygen  volume  taken  it  will  be  seen  that  the 


FIG.  11. 

oxygen  actually  used  is  just  one-half  of  the  hydrogen 
volume  taken.  The  droplet  of  water  formed  in  the  ex- 
plosion is  so  small  that  its  volume  may  be  neglected  in 
comparison  with  the  gas  volumes  concerned. 

In  measuring  the  gases  in  this  experiment  certain  pre- 
cautions must  be  taken  which  are  fully  explained  in  a  fol- 
lowing chapter.  In  all  such  cases  temperature  and 
pressure  must  be  accurately  observed  so  that  by  reduction 
all  the  volumes  may  be  compared  under  the  same  condi- 
tions. 


GENERAL  CHEMISTRY.  59 

Physical  Properties  of  Water.  Pure  water  has  a 
constant  freezing  point  and  constant  boiling  point  under 
constant  conditions  of  pressure.  It  is  a  poor  conductor 
of  heat  and  is  practically  a  nonconductor  of  electricity. 
When  water  at  the  ordinary  temperature  is  heated  it 
expands  pretty  regularly  until  its  boiling  point  is  reached, 
and  by  further  addition  of  heat  is  converted  into  steam, 
the  volume  of  which  is  about  1,700  times  that  of  the  water. 
When  cooled  a  contraction  of  volume  follows  until  the 
temperature  of  4°  C.  is  reached.  At  this  temperature,  or 
more  accurately,  3.95°,  water  reaches  its  smallest  volume 
for  a  given  weight,  and  therefore,  its  maximum  density. 
When  further  cooled  it  expands  slightly  until  the  freezing 
point  is  reached.  On  conversion  into  ice  a  marked  expan- 
sion takes  place,  100  volumes  of  water  at  0°  yielding  109.1 
volumes  of  ice. 


Capacity  for  Heat.  Unit  of  Heat.  It  is  a  matter  of 
common  experienc£~that  water  "absorbs"  a  large  amount 
of  heat,  practical  application  of  which  power  is  made  in 
the  hot  water  system  of  heating  houses.  In  order  to 
measure  any  amount  of  heat  under  consideration  phys- 
icists have  adopted  what  is  known  as  a  unit  of  heat.  This 
may  be  arbitrarily  defined  as  the  amount  of  heat  necessary 
to  raise  the  temperature  of  a  gram  of  pure  water  through  one 
centigrade  degree.  To  be  scientifically  correct  this  defini- 
tion requires  a  slight  modification  or  qualification,  but  it  is 
sufficiently  close  for  the  present  purpose.  Seme  interest- 
ing facts  have  been  brought  out  in  studying  the  absorption 
of  heat  by  water  in  its  three  forms.  To  illustrate  these  let 
us  start  with  a  gram  of  ice  at  0°  centigrade,  that  is,  at  its 
melting  point.  To  convert  this  into  water  at  the  same 
temperature,  a  relatively  large  amount  of  heat  must  be 
applied.  It  has  been  found  that  79.5  units  must  be  added 
to  accomplish  this.  In  other  words,  as  much  heat  must  be 
absorbed  by  the  ice  in  melting  as  would  be  sufficient  to 
warm  79.5  grams  of  water  one  degree,  or  one  gram  of 
water  79.5  degrees.  This  absorbed  heat  is  usually  spoken 
of  as  latent  heat,  because  it  is  lost  or  hidden  as  far  as  any 


60  GENERAL  CHEMISTRY. 

thermometric  observation  is  concerned.  It  is  used  up, 
however,  in  doing  a  certain  kind  of  work  on  the  ice,  that  is 
in  changing  its  physical  condition.  If  we  continue  the 
addition  of  heat  after  the  ice  has  been  melted  the  effect 
now  becomes  visible  in  the  increase  of  the  temperature  of 
the  water.  For  each  unit  of  heat  added  the  temperature  of 
the  water  rises  one  degree  centigrade.  Finally,  on  addi- 
tion of  100  units  of  heat  the  water  begins  to  boil,  and  a 
centigrade  thermometer  immersed  in  it  marks  100  degrees. 
Supposing  the  water  in  a  vessel  under  constant  pressure, 
that  of  the  atmosphere,  it  is  now  observed  that  further 
addition  of  heat  produces  no  corresponding  elevation  of 
temperature.  It  is  a  well-known  fact  that  a  thermometer 
immersed  in  pure  boiling  water  registers  a  constant  tem- 
perature. The  heat  applied  is,  therefore,  again  rendered 
latent  as  in  the  case  of  the  conversion  of  ice  into  water. 

Now  it  is  used  up  in  doing  a  new  kind  of  work,  the 
conversion  of  the  water  into  steam  at  the  same  tempera- 
ture. Accurate  experiments  have  shown  that  about  536 
units  of  heat  are  required  to  convert  a  gram  of  water  at  a 
temperature  of  100°  into  steam  at  the  same  temperature. 
When  a  gram  of  steam  at  this  temperature  condenses  it 
gives  out  the  536  units.  Many  of  thetechnical  applications 
of  water  and  steam  depend  on  these  remarkable  properties. 
Many  other  liquid  and  solid  substances  may  be  brought  to 
a  higher  temperature  than  is  possible  with  water,  but  for  a 
given  range  of  temperature  not  one  of  them  absorbs  nearly 
as  much  heat.  A  hot  water  bag  filled  with  water  at  60°  C. 
will  give  out,  as  it  cools  down,  nearly  ten  times  as  much 
heat  as  would  a  mass  of  iron  of  the  same  weight  at  the 
same  temperature.  The  specific  heat  of  iron  and  all  other 
solid  and  liquid  substances  is  much  below  that  of  water. 
This  matter  will  be  taken  up  again. 

Boiling  Point  of  Water  and  Vapor  Tension.    By 

definition  water  is  said  to  boil  at  a  temperature  of  100° 
centigrade.  But  this  temperature  of  ebullition  depends 
on  the  pressure  on  the  surface  of  the  water,  usually  that  of 
the  air.  If  the  atmospheric  pressure  is  normal,  that  is,  if  it 
is  equivalent  to  the  pressure  of  a  column  of  mercury  760 


GENERAL  CHEMISTRY. 


61 


Mm.  in  height  the  boiling  point  is  constant  at  100°.  At 
pressures  below  this  the  boiling  point  is  below  100°,  and 
at  higher  pressures  it  is  above.  In  any  case  water  boils 
when  the  tension  of  the  vapor  which  it  gives  off  is  equal  to 
the  pressure  of  the  atmosphere. 

When  water  is  heated  in  a  confined  space,  as  in  a  boiler 
furnished  with  a  weighted  valve,  the  temperature  of  ebulli- 
tion may  become  very  high.  Corresponding  to  any  given 
pressure  on  the  surface  of  the  water  there  is  a  definite 
temperature  of  ebullition.  Even  at  very  low  temperatures 
the  vapor  given  off  from  water  possesses  a  certain  tension. 
This  is  shown  in  the  table  below  where  the  tensions  cor- 
responding to  certain  high  as  well  as  low  pressures  are 
given.  The  tensions  are  expressed  in  millimeters  of  mer- 
cury. 


Temp. 
Deg. 

Tension. 

Temp. 
Deg. 

Tension. 

Temp 
Deg. 

Tension. 

Temp. 
Deg. 

Tension. 

1 

4.909 

18 

15.330 

35 

41.78 

120 

1491 

2 

5.272 

19 

16.319 

40 

54.87 

125 

1744 

3 

5.658 

20 

17.363 

45 

71.36 

130 

2030 

4 

6.069 

21 

18.466 

50 

91.98 

135 

2354 

5 

6.507 

22 

19.630 

55 

117.52 

140 

2718 

6 

C.672 

23 

20.858 

60 

148.88 

145 

3126 

7 

7.466 

24 

22.152 

65 

187.10 

150 

3581 

8 

7.991 

25 

23.517 

70 

233.31 

155 

4089 

9 

8.548 

26 

24.956 

75 

288.76 

160 

4652 

10 

9.140 

27 

26.471 

80 

354.87 

165 

5275 

11 

9.767 

28 

28.065 

85 

433.19 

170 

5962 

12 

10.432 

29 

29.744 

90 

525.47 

175 

6717 

13 

11.137 

30 

31.510 

95 

633.66 

180 

7546 

14 

11.884 

31 

33.37 

100 

760.00 

185 

8453 

15 

12.674 

32 

35.32 

105 

906. 

190 

9443 

16 

13.510 

33 

37.37 

110 

1075. 

195 

10510 

17 

14.395 

34 

39.52 

115 

1269. 

200 

11689 

From  the  table  it  is  seen  that  at  a  temperature  of  150° 
the  tension  of  the  vapor  on  the  containing  vessel  is  3,581 
millimeters  of  mercury,  or  over  4.7  atmospheres. 


Water  as  a  Standard  Substance.     Because  of  the 
fact  that  water  is  everywhere  abundant  and  readily  ob- 


62  GENERAL  CHEMISTRY. 

tained  in  pure  condition,  it  is  well  suited  for  use  as  a 
standard  substance.  It  is  employed  in  the  definition  of  a 
unit  of  heat  as  mentioned  above.  In  the  construction  of 
thermometer  scales  it  is  employed  in  determining  two 
fixed  points.  In  the  centigrade  or  Celsius  scale  one  fixed 
point  is  called  the  zero  point,  and  represents  the  tempera- 
ture at  which  water  freezes.  The  other  fixed  point  is 
marked  100°  and  represents  the  boiling  point  of  water  at 
the  normal  pressure  of  the  air  at  sea  level.  By  definition 
a  gram  is  taken  as  the  weight  of  a  cubic  centimeter  of  pure 
water  at  a  temperature  of  4°  C.  at  the  latitude  of  Paris. 
By  specific  gravity  or  specific  weight  we  understand  the 
ratio  of  the  weight  of  a  given  volume  of  substance  to  the 
weight  of  the  same  volume  of  pure  water  as  the  standard. 
In  many  other  cases  water  is  taken  as  one  of  the  elements 
in  comparison,  but  the  above  illustrations  are  sufficient. 

Solvent  Action  of  Water.  Water  is  the  best  single 
solvent  known,  although  for  practical  purposes  many  sub- 
stances may  be  considered  as  insoluble  in  it.  Of  those 
which  do  dissolve  some  are  much  more  soluble  than 
others.  One  volume  of  water  dissolves  0.041  volume  of 
oxygen  at  a  temperature  of  0°  C.  and  a  pressure  of 
760  Mm.,  but  under  the  same  conditions  it  dissolves 
nearly  80  volumes  of  sulphurous  oxide  and  about  500  vol- 
umes of  hydrochloric  acid  gas.  Some  mineral  substances, 
as  gypsum,  are  but  slightly  soluble  while  others,  salt  and 
saltpeter  for  illustrations,  dissolve  very  largely.  With  in- 
crease of  temperature  there  is  in  most  cases  a  marked 
increase  in  the  degree  of  solubility  of  bodies.  Even  such 
hard  substances  as  the  glass  of  our  beakers  and  test-tubes 
dissolve  to  a  very  slight  extent  by  long  contact  with  boiling 
water.  This  fact  is  of  great  importance  in  some  branches 
of  chemical  analysis. 

Natural  Waters.  All  natural  water  comes  to  the 
earth  in  the  form  of  rain,  and  in  this  condition  it  is  nearly 
pure,  containing  not  much  more  than  traces  of  dissolved 
gases.  This  is  especially  true  of  the  water  collected  at 
the  end  of  a  shower,  that  which  falls  first  serving  to 


GENERAL  CHEMISTRY.  63 

down  the  dust  suspended  in  the  air.  On  reaching  the 
earth  the  character  of  the  water  is  very  speedily  modified 
by  the  mineral  substances  which,  by  virtue  of  its  solvent 
action,  it  takes  up.  As  the  rain  descends  it  dissolves  a 
little  carbon  dioxide  and  thus  becomes  a  very  weak  solu- 
tion of  carbonic  acid.  This  acid  aids  in  the  solution  of 
limestone  and  other  substances  from  the  soil,  and  in  this 
way  the  water  becomes  hard.  Hard  waters  are  those 
which  hold  in  solution  relatively  large  amounts  of  certain 
mineral  substances,  principally  salts  of  calcium  and  mag- 
nesium. If  the  soil  on  which  the  rain  falls,  and  through 
which  it  filters  or  percolates,  is  free  practically  from  these 
soluble  mineral  constituents,  the  water  appearing  later 
as  a  spring  or  brook  is  called  soft  water. 

Some  natural  waters  contain  not  over  50  milligrams  of 
dissolved  substances  in  one  liter,  while  others  contain  in 
a  liter  many  grams. 

Purification  of  Waters.  Natural  waters  contain 
substances  dissolved  and  substances  in  suspension.  Both 
may  be  objectionable  for  certain  purposes,  and  before  use 
water  must  often  be  freed  from  them.  The  highest  degree 
of  purification  is  ordinarily  accomplished  by  distillation, 
the  condensed  steam  being  free  from  the  dissolved  solids 
and  suspended  matters  originally  present.  In  many  cases 
purification  extends  only  to  a  mechanical  separation  of 
suspended  matters,  which  is  accomplished  by  filtration 
through  sand,  charcoal  or  other  porous  substance.  Fi- 
nally, water  is  frequently  treated  on  the  large  scale  with 
something  that  produces  in  it  a  bulky  precipitate  which  in 
settling  carries  down  practically  everything  in  suspension. 
After  such  precipitation,  by  milk  of  lime  or  alum,  the  water 
is  allowed  to  settle  thoroughly,  or  clarification  is  hastened 
by  filtration. 

HYDROGEN  DIOXIDE. 

Hydrogen  and  oxygen  combine  in  two  proportions.  In 
one  case  water  is  the  result  and  in  the  other  the  body 
known  as  hydrogen  dioxide  or  peroxide. 


64  GENERAL  C HEM  IS  TR  Y. 

Occurrence.  This  substance  occurs  in  some  atmos- 
pheres in  small  traces,  being  produced  by  several  natural 
agencies.  In  the  air  it  is  frequently  confounded  with 
ozone,  as  the  two  bodies  are  much  alike  in  their  behavior 
with  reagents. 

Preparation.  Hydrogen  dioxide  can  easily  be  made 
in  the  laboratory  as  illustrated  by  the  next  experiment.  Its 
preparation  depends  on  the  decomposition  of  a  solid  sub- 
stance, usually  barium  dioxide,  by  means  of  sulphuric 
acid. 

Ex.  58.  Take  about  10  grams  of  barium  dioxide  in  a  beaker, 
moisten  it  with  water  and  allow  the  mixture  to  stand  half  an  hour  or 
longer,  with  occasional  stirring.  Then  add  about  20  cubic  centimeters 
of  dilute  sulphuric  acid  (1  vol.  of  acid  to  10  of  water),  stir  well  and  af- 
ter a  few  minutes  filter.  The  acid  acts  on  the  barium  dioxide  in  the 
dry  state  very  slowly,  but  if  the  latter  has  been  previously  hydrated  by 
the  action  of  water  the  decomposition  is  much  more  rapid.  Insoluble 
barium  sulphate  and  soluble  liquid  hydrogen  dioxide  result.  The  former 
is  separated  by  filtration  while  the  solution  of  the  latter  serves  for 
tests. 

Properties.  As  usually  made,  hydrogen  dioxide  is 
largely  mixed  with  water,  but  in  the  pure  state  it  is  a'thick- 
ish  liquid  with  a  specific  gravity  of  1.45.  In  this  condi- 
tion it  is  not  stable,  but  decomposes  at  the  ordinary 
temperature  into  water  and  oxygen  gas.  For  this  reason 
the  substance  is  always  kept  in  very  dilute  condition,  and 
experience  has  shown  that  its  stability  is  increased  by 
having  a  little  free  sulphuric  acid  in  the  solution. 

Even  in  dilute  solution  it  is  characterized  by  its  strong 
oxidizing  properties,  and  the  numerous  uses  it  has  found 
in  the  arts  and  medicine  depend  on  this  fact. 

Tests.  Some  of  the  reactions  of  hydrogen  dioxide 
may  be  shown  by  these  experiments: 

Ex.  59.  To  a  solution  of  the  dioxide  add  some  solution  of  potas- 
sium iodide.  A  decomposition  of  this  compound  takes  place,  with 
liberation  of  iodine,  as  is  shown  by  the  brown  color  of  the  liquid.  If  a 
little  starch  paste  is  added,  it  turns  blue.  The  solution  of  the  dioxide 
used  in  this  test  should  be  nearly  neutral,  that  is,  free  from  anything 
greater  than  traces  of  acid.  This  test,  it  will  be  observed,  is  very  simi- 
lar to  that  for  ozone,  given  earlier. 


GENERAL  CHEMISTRY.  65 

Ex.  60.  Another  very  interesting  decomposition  is  shown  as  fol- 
lows: Acidify  the  solution  of  the  dioxide  with  dilute  sulphuric  acid  and 
add  to  it  a  dilute  aqueous  solution  of  potassium  permanganate,  a  few 
drops  at  a  time.  This  latter  solution  has  a  deep  purple  color  and  as  the 
drops  fall  into  the  dioxide  liquid  the  color  disappears,  while  bubbles  of 
oxygen  gas  escape.  On  addition  of  a  sufficient  amount  of  the  perman- 
ganate solution  the  purple  color  persists.  The  chemical  reaction  taking 
place  here  is  somewhat  complicated  and  cannot  be  explained  in  this 
stage  of  the  work. 

The  two  experiments  just  given  illustrate  the  marked 
property  of  producing  decompositions  possessed  by  the 
dioxide.  This  property  depends  on  the  fact,  to  be  more 
fully  explained  later,  that  part  of  the  oxygen  united  to 
hydrogen  in  the  compound  is  only  loosely  held.  It  is  very 
easily  liberated  and  performs  highly  characteristic  oxida- 
tion changes  in  consequence. 

Hydrogen  dioxide  is  used  to  some  extent  in  laboratories 
as  a  reagent  but  its  most  common  applications  are  as  a 
bleaching  agent  and  in  medicine. 


CHAPTER   III. 


CHLORINE  AND  HYDROCHLORIC  ACID THEO- 
RETICAL CONSIDERATIONS. 

WE  COME    now    to    the   consideration    of    some    very 
important  substances  which  never  occur  in  the  free 
state    in    nature,   but    in    many    compounds    are    widely 
distributed. 

CHLORINE. 

Occurrence.  This  is  an  exceedingly  abundant  sub- 
stance in  combination,  being  one  of  the  constituents  of 
common  salt.  The  other  constituent  is  a  body  called 
sodium.  It  is  found  also  in  many  compounds  somewhat 
similar  to  salt,  all  of  which  are  called  chlorides.  Salt  is 
known  as  sodium  chloride,  while  the  next  most  abundant 
chloride  is  potassium  chloride. 

History.  Chlorine  was  first  prepared  by  the  great 
Swedish  chemist,  Scheele,  in  1774,  and  by  a  method  which 
is  still  commonly  employed  for  the  purpose,  viz.,  by  the 
action  of  black  oxide  of  manganese  on  hydrochloric  acid. 
Scheele  did  not,  however,  recognize  the  true  nature  of  the 
substance,  and  it  remained  for  Humphrey  Davy,  in  1810, 
to  supply  this  information  and  propose  the  name  now  given 
to  the  body. 

Preparation.  All  methods  of  preparing  chlorine  de- 
pend on  the  decomposition  of  some  of  the  chlorides.  Usu- 
ally we  employ  sodium  chloride  or  hydrogen  chloride, 
called  also  hydrochloric  acid.  The  preparation  by  the  use 
of  hydrochloric  acid  will  be  illustrated  first.  This  sub- 


GENERAL  CHEMISTRY. 


67 


stance,  as  will  appear  later,  is  a  compound  of  hydrogen 
and  chlorine,  and  our  problem  is  to  separate  one  from  the 
other.  The  next  experiment  will  show  how  this  can  be 
done. 

Ex.  61.  In  a  flask  holding  300  Cc.  or  more,  take  about  50  Gm.  of 
manganese  dioxide,  the  substance  already  used  with  potassium  chlorate 
in  the  preparation  of  oxygen.  Pour  over  the  dioxide  about  200  Cc.  of 
commercial  strong  hydrochloric  acid.  Close  the  flask  with  a  stopper 


FIG.   12. 


having  two  perforations.  Through  one  of  these  a  funnel  tube  passes, 
the  lower  end  dipping  into  the  acid.  A  delivery  tube  passes  out  through 
the  other  perforation.  This  extends  up  about  6  or  8  Cm.,  and  there  is 
bent  at  right  angles,  the  horizontal  portion  following  having  about  the 
same  length.  To  this,  by  means  of  a  rubber  connection,  is  joined  a 
second  bent  glass  tube,  the  longer  limb  of  which  has  a  length  of  about 
20  Cm.  The  flask  is  supported  on  a  sand-bath,  as  shown  in  the  above 
figure,  while  the  delivery  tube  extends  down  into  an  empty  dry  bottle  of 
200  to  300  Cc.  capacity.  When  the  flask  is  charged  and  properly 
mounted  apply  a  gentle  heat  to  the  sand-bath.  This  hastens  the  action 
of  the  acid  on  the  manganese  dioxide.  A  greenish  yellow  gas  soon  fills 


68  GENERAL  CHEMISTRY. 

the  flask  and  passes  over  into  the  dry  bottle.  The  delivery  tube  should 
enter  the  bottle  through  a  piece  of  perforated  cardboard,  in  order  to 
prevent,  as  far  as  possible,  the  escape  of  the  gas  while  the  bottle  is  fill- 
ing. As  chlorine  gas  is  much  heavier  than  the  air,  it  can  be  collected 
in  this  manner  by  displacement,  the  air  being  driven  up  out  of  the 
bottle.  When  the  bottle  appears  to  be  quite  full  of  the  gas  remove  it 
and  put  a  second  in  its  place.  Collect  four  or  five  bottles  of  the  gas. 
Each  bottle  as  removed  from  the  generator  must  be  covered  by  a  glass 
plate.  After  collecting  the  desired  amount  of  the  gas,  replace  the  dry 
bottles  by  one  containing  water.  Continue  the  application  of  heat  and 
add  more  hydrochloric  acid  if  necessary.  Chlorine  gas  is  somewhat 
soluble  in  water,  and  in  this  manner  a  solution  is  obtained  which  is  used 
in  a  following  experiment.  All  experiments  with  chlorine  must  be  per- 
formed in  a  fume  closet.  Save  the  liquid  in  the  generator. 

We  have  now  several  bottles  of  the  gas,  of  which  the 
greenish  yellow  color  is  characteristic.  The  odor  also  is 
marked  and  disagreeable  in  the  extreme.  The  student 
must  avoid,  as  far  as  possible,  inhaling  it,  as  it  is  exceed- 
ingly irritating  to  the  air  passages.  In  large  amounts  it  is 
even  very  dangerous.  For  this  reason  the  direction  is 
given  to  carry  out  all  experiments  in  a  fume  closet  where 
there  is  sufficient  draught  to  carry  off  any  escaping  gas. 
Having  become  aware  of  the  more  prominent  features  of 
the  substance,  the  student  may  determine  some  of  its  prop- 
erties by  experiment. 

Ex.  62.  Chlorine  is  a  strong  bleaching  agent,  which  maybe  shown 
as  follows:  Moisten  a  strip  of  colored  calico  in  water  and  suspend  it  in 
one  of  the  bottles  of  the  gas,  replacing  the  glass  plate  after  introducing 
the  fabric.  In  time  the  color  fades  through  the  destructive  action  of 
the  chlorine.  In  the  dry  state  the  gas  is  practically  without  action, 
hence  the  direction  given  to  moisten  the  calico.  In  the  same  bottle 
pour  a  little  solution  of  indigo.  If  the  bleaching  of  the  calico  has  not 
removed  all  the  chlorine  the  indigo  color  will  be  destroyed  likewise. 

This  important  property  of  the  gas  is  utilized  on  the 

large  scale  in  the  wholesale   bleaching  of  many  articles. 

The  action   however  is  often  destructive  of  the  organic 
fiber  itself. 

Ex.  63.  Chlorine  has  a  marked  affinity  for  many  metals  as  well  as 
for  organic  colors,  and  this  may  be  illustrated  by  a  very  simple  test. 
Remove  the  glass  plate  from  one  of  the  bottles,  filled  as  above  described, 
and  put  in  its  place  a  piece  of  wire  gauze.  Through  this  sift  into  the 
gas  a  little  powdered  antimony.  The  fine  metal  particles  ignite  as  they 
fall  in  a  shower  through  the  gas  and  burn  brightly,  forming  a  chloride 


GENERAL  CHEMISTRY.  69 

of  antimony.  Many  other  metals  burn  equally  as  well  if  they  are  taken 
in  the  form  of  fine  wire  or  foil.  Antimony  in  powder  exposes  a  large 
surface  to  the  action  of  the  gas. 

Chlorine  was  prepared  above  by  decomposing  its 
compound  with  hydrogen.  It  may  readily  be  made  to  com- 
bine with  hydrogen,  again  reproducing  the  acid.  If  equal 
volumes  of  the  two  gases  be  mixed  in  the  dark  and  then 
exposed  to  the  light,  combustion  of  ten  follows  with  a  sharp 
explosion.  The  student  is  not  advised  to  attempt  this  experi- 
ment as  it  is  dangerous,  unless  carried  out  with  certain  pre- 
cautions which  need  not  be  described  here.  The  same 
affinity  of  hydrogen  for  chlorine  may  be  shown  in  an- 
other manner  without  risk  as  will  now  be  explained. 
Many  substances  are  known  which  consist  of  hydrogen 
and  carbon  only.  These  bodies  are  called  hydrocar- 
bons. It  has  been  found  by  experiment  that  chlorine 
gas  is  able  to  decompose  a  number  of  these  substances, 
combining  with  the  hydrogen  to  produce  hydrochloric 
acid,  while  the  carbon  is  set  free  as  a  fine  black  powder  or 
soot.  Among  the  hydrocarbons  which  exhibit  this  behav- 
ior, turpentine  oil  serves  our  purpose  best,  and  wifl  be 
employed  therefore  in  our  test. 

Ex.  64.  Pour  some  oil  of  turpentine  in  a  test-tube,  and  in  it  im- 
merse a  strip  of  filter  paper  which  has  been  twisted  in  the  form  of  a 
taper.  After  withdrawing  the  paper  press  against  it  a  second  piece  of 
dry  filter  paper,  in  order  to  absorb  the  excess  of  the  oil  taken  up.  Now 
remove  the  glass  plate  from  a  chlorine  bottle  and  into  it  dip  the  taper, 
which  in  a  few  seconds  darkens  and  finally  burns  with  a  very  sooty  flame, 
sending  up  a  large  volume  of  smoke.  This  consists  of  the  liberated  car- 
bon, while  hydrochloric  acid  vapors  are  formed,  as  can  be  shown  by 
proper  tests.  For  the  success  of  this  experiment  the  chlorine  gas  must 
be  practically  free  from  air,  and  the  paper  must  not  be  wet  but  only 
moist  from  the  action  of  the  turpentine.  , 

The  same  general  behavior  is  shown  by  burning  a  small 
wax  candle,  or  even  a  small  jet  of  illuminating  gas  from  a 
bent  glass  tube,  in  a  bottle  of  the  gas.  Both  the  wax  and 
illuminating  gas  contain  hydrogen  in  combination  with 
carbon.  A  combustion  of  either  of  these  substances,  begun 
in  the  air,  continues  in  chlorine  with  a  very  sooty  flame. 

A  small  portion  of  the  solution  of  chlorine  water  pre- 
pared above  may  be  used  for  a  bleaching  test,  but  the 


70  GENERAL  CHEMISTRY. 

larger  portion  should  be   preserved  in  a  stoppered  bottle, 
kept  in  a  dark  place,  for  several  tests  to  follow. 

Ex.  65.  Prove  that  the  chlorine  water  bleaches  as  does  the  gas, 
by  immersing  a  piece  of  calico  in  it,  or  by  pouring  in  some  solution  of 
indigo  or  litmus. 

The  method  given  above  serves  very  well  for  making 
small  quantities  of  chlorine  for  experimental  purposes,  but 
foi  the  preparation  of  larger  quantities  the  following  proc- 
ess is  much  better.  . 

The  gas  is  generated  in  a  large  flask  arranged  as  before, 
but  from  a  mixture  of  salt,  manganes'e  dioxide  and  sul- 
phuric acid.  Mix  about  equal  weights  of  the  salt  and  the 
dioxide  ;  pour  this  into  the  flask,  and  then  through  the 
funnel  tube  add  gradually  to  the  mixture  about  twice  its 
weight  of  sulphuric  acid,  previously  diluted  with  an  equal 
weight  of  water.  The  flask  is  slowly  heated  on  a  sand-bath, 
and  the  gas  is  given  off  gradually. 

If  the  gas  is  to  be  used  for  making  pure  chlorine  water 
it  should  be  led  through  a  wash  bottle  first.  In  its  sim- 
plest form  this  may  consist  of  a  wide  mouth  bottle  holding 
about  250  Cc.,  closed  with  a  stopper  with  two  openings. 
Through  one  of  these  a  tube  from  the  generator  passes  and 
dips  beneath  the  surface  of  water,  which  about  half  fills  the 
bottle.  A  second  tube  passes  up  from  the  under  surface  of 
the  stopper  and  bends  over,  to  lead  the  gas  to  water  or  to 
collecting  bottles,  as  desired.  By  this  arrangement  the 
gas  is  washed  by  bubbling  through  the  water  in  the  wash 
bottle.  It  loses  small  amounts  of  hydrochloric  acid  and 
other  impurities  carried  over  from  the  generator,  as  these 
are  more  soluble  in  water  than  chlorine  is. 

At  20°  C.  1  volume  of  water  dissolves  2.15  volumes  of 
chlorine,  and  the  solution  so  made  can  be  kept  a  long  time 
in  a  glass  stoppered  bottle  in  the  dark.  But  exposed  to 
light,  gradual  decomposition  takes  place,  the  hydrogen  of 
the  water  combining  with  the  chlorine  to  form  hydrochlo- 
ric acid,  while  the  oxygen  is  liberated.  In  bright  sunlight 
the  action  is  comparatively  rapid,  as  can  very  easily  be 
shown  by  experiment. 

Ex.  66.     Fill  a  liter  bottle  with  distilled  water   and   pass  chlorine 


GENERAL  CHEMISTRY.  71 

gas  into  it  to  complete  saturation.  Close  the  bottle  then  with  a  stopper, 
so  as  to  exclude  all  air,  and  invert  it  in  a  jar  containing  strong  chlorine 
water.  Remove  the  stopper  and  stand  the  jar,  with  inverted  bottle,  in  a 
window  exposed  to  sunlight.  In  a  short  time  gas  bubbles  will  be  seen 
to  ascend  through  the  liquid  in  the  bottle.  With  average  sunlight 
several  days  may  elapse  before  the  reaction  is  complete.  The  gas  col- 
lected over  the  water  in  the  bottle  may  now  be  tested.  Insert  the  stop- 
per in  the  bottle,  bring  it  to  the  upright  position,  then  withdraw  the 
stopper  and  apply  the  usual  test  for  oxygen  by  means  of  a  glowing 
splinter.  The  reaction  is  distinct  and  characteristic.  If  the  bottle  is 
allowed  to  stand  long  enough  in  the  light  the  green  color  and  the  odor 
of  the  chlorine  disappear,  while  the  sharp,  hydrochloric  acid  odor  can 
be  noticed. 

Other  Methods  of  Preparation.  Chlorine  can  be 
liberated  by  several  other  reactions,  some  of  which  have 
technical  importance.  One  of  these  may  be  illustrated 
here  by  a  brief  experiment. 

Ex.  67.  Take  a  gram  or  two  of  powdered  potassium  dichromate  in 
a  test-tube  and  pour  over  it  a  few  cubic  centimeters  of  strong  commer- 
cial hydrochloric  acid.  Apply  heat.  Decomposition  of  the  hydro- 
chloric acid  soon  takes  place  as  shown  by  the  appearance  of  greenish 
yellow  fumes  in  the  test-tube;  that  these  consist  of  chlorine  may  be 
inferred  from  the  color,  odor  and  bleaching  action  easily  determined. 
An  explanation  of  this  reaction  will  be  given  later. 

A  strong  solution  of  hydrochloric  acid  may  be  decom- 
posed by  electricity  in  apparatus  similar  to  that  employed 
for  the  electrolysis  of  water.  Several  applications  are 
made  of  this  fact. 

On  the  large  scale  chlorine  is  liberated  in  quantity  and 
cheaply  by  what  is  known  as  the  Deacon  process,  from  the 
name  of  the  discoverer.  In  this  process  a  stream  of  hydro- 
chloric acid  gas  is  blown  along  with  a  stream  of  air  through 
a  series  of  heated  tubes  containing  crushed  brick  impreg- 
nated with  copper  sulphate.  The  oxygen  of  the  air  takes 
the  hydrogen  of  the  acid  and  chlorine  is  left  free,  mixed 
with  the  nitrogen  of  the  air.  The  crude  chlorine  so  made 
is  suitable  for  the  production  of  bleaching  powder  and 
other  products. 

General  Tests  for  Chlorine.  The  common  proper- 
ties of  chlorine  are  so  marked  that  they  afford  easy  means 
of  recognition.  Like  ozone  and  hydrogen  peroxide,  chlo- 


73  GENERAL  CHEMISI'RY. 

rine  is  able  to  decompose  potassium  iodide,  and  hence  the 
so-called  ozone  test-paper,  described  in  an  earlier  section, 
serves  also  as  a  chlorine  test,  when  applied  in  moist  con- 
dition to  the  gas  supposed  to  contain  or  consist  of  chlo- 
rine. The  methods  by  which  chlorine  may  be  recognized 
when  mixed  with  other  gases  are  described  in  Qualitative 
Analysis. 

Physical  Properties.  As  stated  above,  chlorine  is 
somewhat  soluble  in  water.  One  volume  of  water  at  20° 
dissolves  about  2.15  volumes  of  chlorine.  With  ice  water 
it  forms  a  crystalline  compound.  Under  strong  pressure 
gaseous  chlorine  may  be  condensed  to  a  liquid  having  a 
specific  gravity  of  1.33.  This  liquid  is  now  an  article  of 
commerce.  One  liter  of  chlorine  gas,  at  0°  and  under  a 
pressure  of  760  Mm.,  weighs  3.18  grams. 

Uses  of  Chlorine.  While  chlorine  has  many  applica- 
tions on  the  large  scale,  it  has  also  some  in  the  laboratory. 
The  gas  itself  is  frequently  used  and  also  the  solution,  or 
chlorine  water.  It  was  directed  above  to  save  some  of  this 
solution,  and  with  it  several  experiments  will  be  made,  as 
explained  a  few  pages  in  advance. 


HYDROCHLORIC    ACID. 

History.  This  acid  was  known  in  crude  form  to  the 
Arabian  chemists  and  was  made  from  salt  and  green  vitriol 
in  the  15th  century.  About  the  middle  of  the  17th  century 
it  was  first  made  by  a  process  like  that  employed  to  day. 

It  has  been  already  intimated  that  this  substance  may 
be  formed  by  the  direct  union  of  chlorine  with  hydrogen, 
also  by  the  action  of  chlorine  on  water  in  sunlight.  It  is 
usually  prepared,  however,  by  the  decomposition  of  a 
chloride  by  means  of  sulphuric  acid.  The  cheapest  chloride 
known  is  sodium  chloride,  or  common  salt,  and  hence  this 
substance  is  nearly  always  employed  in  the  preparation. 
The  reaction  may  be  carried  out  very  easily  as  a  laboratory 
experiment,  by  a  method  now  to  be  given. 


GENERAL  CHEMISTRY. 


7:* 


Ex.  68.     Arrange  apparatus  as  shown  by  the  illustration. 

The  flask,  to  the  left,  on  a  sand-bath  has  a  capacity  of  about 
500  Cc.  It  is  charged  with  about  50  grams  of  common  salt,  and  is 
closed  by  a  stopper  with  two  perforations,  through  one  of  which  passes 
a  funnel  tube  leading  nearly  to  the  bottom  of  the  flask.  Through  the 
other  perforation  a  delivery  tube  passes,  and  this  ends  finally  in  a 
Woulfe  bottle,  half  filled  with  water,  but  the  delivery  tube  must  not 
dip  beneath  the  surface  of  the  water  here.  Another  tube  leads  from  this 
first  Woulfe  bottle  into  a  second,  likewise  half  filled  with  water.  In 
this  case  the  tube  dips  beneath  the  surface  of  the  water.  From  the 
second  Woulfe  bottle  a  tube  leads  to  a  flask  of  water.  Each  one  of  the 
Woulfe  bottles  has  three  openings.  Through  one  of  these,  in  each  case, 
a  so-called  safety  tube  passes  and  dips  into  the  water.  The  object  of 
these  safety  tubes  is  to  provide  for  easy  communication  with  the  air  in 
case  the  pressure  of  gas  in  the  generating  flask  should  suddenly  diminish 


FIG.  13. 


and  the  lower  end  of  the  funnel  tube  should  be  closed  by  the  materials 
around  it.  When  the  apparatus  is  in  order,  the  flask  containing  the 
salt,  as  explained,  pour  in  about  50  Cc.  of  strong  sulphuric  acid  through 
the  funnel  tube,  a  little  at  a  time. 

Immediately  a  very  lively  reaction  begins.  The  mass  in  the  flask 
froths  and  rises,  while  gas  bubbles  escape  through  the  water  in  the 
second  Woulfe  bottle  and  the  adjoining  flask.  This  is  air  being 
expelled.  The  hydrochloric  acid  formed  by  the  action  of  the  sulphuric 
acid  is  now  seen  to  enter  the  first  Woulfe  bottle  and  pass  down  from 
the  end  of  the  delivery  tube  to  the  water  and  mix  with  it.  The 
remarkable  affinity  of  the  water  and  gas  is  illustrated  by  this.  After  a 
time  the  flow  of  gas  from  the  generator  lessens  and  then  heat  should  be 
applied  to  the  sand-bath.  At  the  beginning  this  is  not  necessary  as  the 
two  substances  react  on  each  other  in  the  cold.  After  a  time  the  action 


74  GENERAL  CHEMISTRY. 

in  the  flask  ceases  and  no  more  gas  passes  over  into  the  Woulfe  bottles. 
These  are  detached  from  the  generator  and  their  contents  tested. 

Ex.  69.  Begin  the  tests  by  taking  equal  volumes  of  liquid,  as  a 
test-tube  full,  from  the  two  Woulfe  bottles  and  the  small  flask.  Pour 
the  contents  of  the  test-tubes  into  three  small  clean  beakers,  and  add 
to  each  a  few  drops  of  solution  of  silver  nitrate.  In  the  liquid  from  the 
first  Woulfe  bottle  a  heavy,  curdy  white  precipitate  forms,  in  that  from 
the  second  the  amount  of  precipitate  is  much  less,  while  in  the  case  of 
the  liquid  from  the  end  flask  an  opalescence  only  may  result. 

Ex.  70.  Remove  now  three  more  equal  portions  as  before,  trans- 
fer them  to  clean  beakers  and  add  to  each  five  drops  of  a  weak  alcoholic 
solution  of  phenol-phthalein,  which  may  cause  a  faint  opalescence  of  no 
consequence.  Have  at  hand  a  dilute  solution  of  sodium  hydroxide 
(caustic  soda)  and  add  this  gradually  to  the  contents  of  each  beaker, 
beginning  with  that  from  the  small  flask.  One  or  two  drops  of  sodium 
hydroxide  solution  may  be  sufficient  to  impart  a  red  color  to  the  liquid, 
and  this  indicates  that  the  acid  present  has  been  fully  neutralized  by 
the  solution  added,  which  is  an  alkali.  The  phenol-phthalein  is  a  sub- 
stance which  is  turned  bright  red  by  alkaline  solutions  and  hence  it  is 
employed  here  to  show  a  change  from  acid  to  alkaline  condition.  It  will 
be  found  that  to  neutralize  the  acid  from  the  second  Woulfe  bottle 
more  of  the  soda  solution  must  be  used,  while  for  the  liquid  from  the 
first  Woulfe  bottle  a  very  large  volume,  relatively,  of  the  alkali  solution 
is  necessary.  It  appears  from  this  test  that  in  the  experiment  most  of 
the  hydrochloric  acid  generated  remains  in  the  first  Woulfe  bottle. 

The  last  two  experiments  show  certain  important  prop- 
erties of  hydrochloric  acid.  It  gives  a  precipitate  with 
solution  of  silver  nitrate,  it  neutralizes  a  strong  alkali 
solution,  and  it  is  very  soluble  in  water.  We  know  this 
last  to  be  true  because  both  the  silver  and  the  alkali  tests 
show  that  the  most  of  the  acid  is  in  the  water  of  the  first 
bottle.  That  the  acid  is  a  gaseous  body  is  indicated  by 
the  manner  in  which  it  is  liberated  in  the  beginning  of  the 
process  when  the  sulphuric  acid  is  first  added  to  the  salt. 
Without  application  of  heat,  it  was  seen  that  something 
passed  over  from  the  generating  flask  into  the  collecting 
bottle  and  this  was  evidently  a  gas. 

It  has  been  found  by  experiment  that  at  0°  C.  1  vol- 
ume of  water  absorbs  very  nearly  500  volumes  of  the  gas. 
The  common  hydrochloric  acid  which  we  handle  in  liquid 
form  is  merely  an  aqueous  solution  of  the  gas,  containing 
from  25  to  40  per  cent  by  weight  of  the  real  acid.  For 
many  purposes  the  strong  acid  is  diluted,  before  use,  with 


GENERAL  CHEMISTRY.  75 

more  water.  With  about  10  cubic  centimeters  of  the 
strong  acid  solution  in  the  first  Woulfe  bottle  make  the 
following  experiment. 

Ex.  71.  Transfer  the  acid  to  a  small  heaker  and  add  a  few  drops 
of  the  phenol-phthalein  indicator.  Then  pour  in  very  gradually,  as 
before,  some  caustic  soda  solution  until  a  red  color  just  appears.  Now, 
by  means  of  a  glass  rod,  add  a  drop  or  two  more  of  the  acid,  or  sufficient 
to  discharge  the  color.  This  yields  a  very  nearly  neutral  solution,  the  acid 
being  only  slightly  in  excess.  Pour  the  liquid  into  a  clean  porcelain 
evaporating  dish  and  boil  it  down  to  dryness.  If  during  the  evaporation 
the  red  color  returns  it  shows  that  insufficient  acid  has  been  added  and 
a  drop  or  two  more  may  be  mixed  with  the  liquid.  When  the  evapora- 
tion is  complete  heat  strongly  a  few  minutes  longer,  allow  to  cool  and 
observe  the  taste.  We  have  here  common  salt,  similar  to  that  decom- 
posed in  the  large  flask  in  the  reaction  with  sulphuric  acid. 

In  this  series  of  experiments  we  have  illustrations  of 
some  very  important  chemical  reactions.  In  the  making 
of  hydrochloric  acid  we  decompose  salt,  which  is  a  com- 
pound body,  and  secure  two  new  substances.  One  of  these 
is  the  gas  hydrochloric  acid  which  distills  over,  while  the 
other  is  a  solid  substance,  left  in  the  generating  flask, 
and  is  known  as  sodium  sulphate.  When  it  is  seen  that 
in  neutralizing  the  hydrochloric  acid  with  the  soda  solution 
we  reproduce  salt  it  becomes  evident  that  the  soda  adds  in 
the  last  experiment  that  which  the  sulphuric  acid  must 
have  separated  in  the  first.  Hydrochloric  acid  on  the  large 
scale  is  often  produced  as  a  by-product  in  the  manufacture 
of  alkali  from  sodium  or  potassium  chloride,  as  will  be 
explained  later.  Sodium  sulphate  and  potassium  sul- 
phate, formed  as  above  illustrated,  are  converted  into 
sodium  and  potassium  carbonates. 

Experiment  shows  that  this  acid  is  very  active  in  the 
solution  or  decomposition  of  many  bodies.  It  dissolves 
iron,  zinc  and  several  other  metals  forming  chlorides,  its 
hydrogen  being  liberated.  It  dissolves  marble  and  other 
carbonates  with  liberation  of  carbonic  acid  gas  and  forma- 
tion of  chlorides. 

Physical  Properties.  Hydrochloric  acid  can  be  con- 
densed to  a  liquid  under  considerable  pressure  and  in  this 
form  has  a  specific  gravity  of  1.27.  A  liter  of  the  gas, 


76  GENERAL  CHEMISTRY. 

under  standard  conditions,  weighs  1.643  grams.  As  men- 
tioned, the  gas  is  extremely  soluble  in  water,  it  being  pos- 
sible to  prepare  at  a  low  temperature  a  solution  which 
contains  45  per  cent  by  weight  of  the  acid.  This  solution 
is  not  stable  at  higher  temperatures.  A  solution  of  42  per 
cent  strength  can  be  made  to  keep.  Our  strongest  com- 
mercial acid  has  usually  a  strength  of  about  40  per  cent. 
The  stronger  grades  fume  when  exposed  to  the  air  because 
of  a  combination  of  the  real  acid  with  the  moisture  present. 

Uses.  Hydrochloric  acid  is  employed  for  numerous 
purposes  in  the  chemical  laboratory  and  in  technical  oper- 
ations on  the  large  scale.  Much  of  the  crude  by-product, 
mentioned  above,  is  employed  in  the  manufacture  of 
bleaching  powder. 

ELEMENTS  AND  COMPOUNDS. 

Enough  work  has  been  done  thus  far  by  the  student 
to  make  him  acquainted  with  certain  fundamental  differ- 
ences between  bodies.  Oxygen  has  been  prepared  by  sev- 
eral methods,  and  it  was  found  that  it  could  be  readily 
combined  with  several  other  substances.  Hydrogen  and 
chlorine  likewise,  were  secured  in  the  free  pure  condition 
and  in  turn  were  united  with  other  bodies  to  form  new 
substances.  Nowhere  has  anything  been  said  about  the 
decomposition  or  breaking  up  of  oxygen,  hydrogen  and 
chlorine  themselves.  The  question  might  naturally  occur 
to  the  student,  why,  in  experimenting  with  these  three 
bodies,  has  no  experiment  been  given  in  which  they  in 
turn  should  be  decomposed. 

The  fact  is  that  up  to  the  present  time  no  means  have 
been  found  by  which  these  three  substances,  and  many 
others  to  be  mentioned  later,  can  be  resolved  into  any- 
thing simpler.  The  numerous  and  powerful  methods  of 
decomposition  known  to  chemists  have  been  applied  in 
vain  to  the  splitting  of  these  bodies,  and  hence  they  have 
come  to  be  regarded  as  the  real  elements  of  the  material 
world.  They  cannot  be  decomposed,  it  appears,  but  they 
can  combine  to  form  other  substances.  The  new  sub- 


GENERAL  CHEMISTRY.  77 

stances  are  called  compounds,  and  of  these  we  have  al- 
ready had  in  our  work  many  illustrations.  Water,  common 
salt,  hydrochloric  acid,  sulphuric  acid,  mercuric  oxide,  po- 
tassium chlorate,  carbon  dioxide,  hydrogen  dioxide  and 
other  bodies  produced  or  used  were  shown  by  experiment 
to  be  compound  in  their  nature.  In  all  these  cases  the 
existence  of  at  least  two  elements  was  shown  or  could  be 
inferred  with  certainty  from  the  experimental  results. 

There  appear  to  be  about  seventy-four  of  these  ele- 
mentary substances,  while  the  number  of  compound  bodies 
known  is  enormously  large.  That  hydrogen,  oxygen, 
chlorine  and  the  other  so-called  elements  are  really  ele- 
mentary, that  they  can  never  be  decomposed,  yielding  other 
substances,  we  cannot  safely  affirm  ;  indeed,  water  and 
many  other  compound  bodies  were  once  looked  upon  as 
elements.  But  this  much  may  be  safely  said,  that  with  the 
means  now  at  our  command,  we  cannot  decompose  them, 
and,  therefore,  for  all  practical  purposes  they  must  be 
looked  upon  as  elementary. 

ATOMS  AND  MOLECULES. 

A  systematic  or  scientific  study  of  chemical  phenomena 
began  toward  the  end  of  the  last  century,  and,  as  already 
pointed  out,  the  conditions  under  which  many  of  the  ele- 
ments combine  were  soon  recognized.  It  was  found, 
among  other  things,  that  the  power  of  combination  is 
limited;  in  other  words  that  the  elements  can  be  made  to 
unite,  as  a  rule,  in  certain  proportions  only.  Several  cases 
were  known  in  which  metals  combined  with  oxygen  in 
more  than  one  proportion,  but  even  the  crude  analyses  of 
the  time  were  sufficient  to  show  that  in  the  more  highly 
oxidized  bodies  the  amount  of  oxygen  present  is  a  multi- 
ple of  that  in  the  lower.  In  the  two  oxides  of  nitrogen 
known  at  the  beginning  of  this  century,  it  was  found  that 
one  contains  just  twice  as  much  oxygen  as  the  second. 
The  same  relation  was  pointed  out  for  the  two  compounds 
of  carbon  and  oxygen  known,  and  John  Dalton  found  that 
in  two  compounds  of  carbon  and  hydrogen  known  the  ratio 
of  the  weight  of  the  carbon  to  that  of  the  hydrogen  is  just 


78  GENERAL  CHEMISTRY. 

twice  as  great  in  one  case  as  in  the  other.  These  observed 
facts  naturally  caused  much  speculation  among  chemists, 
butDalton  was  the  first  to  propose  a  satisfactory  hypoth- 
esis to  account  for  them. 

From  the  earliest  times  philosophers  were  familiar  with 
the  idea  that  matter  exists  ultimately  in  the  form  of  minute 
indivisible  particles  called  atoms,  and  although  this  view 
was  not  regarded  in  general  as  of  fruitful  importance, 
Dalton  was  able  to  develop  it  further  and  make  finally 
much  of  it.  After  applying  the  conception  to  the  explana- 
tion of  several  purely  physical  phenomena,  he  employed  it 
to  account  for  the  formation  of  chemical  compounds  by  the 
union  of  minute  particles  or  atoms  of  constant  weight. 
According  to  him  an  atom  of  hydrogen  unites  with  an  atom 
of  oxygen  to  form  water;  an  atom  of  hydrogen  with  an 
atom  of  nitrogen  to  form  ammonia;  an  atom  of  hydrogen 
with  an  atom  of  carbon  to  form  ethylene.  A  contemporary 
of  Dalton,  Thomson,  explaining  the  views  of  the  latter, 
mses  this  language: 

"  One  atom  of  a  body,  a,  unites  with  one  atom  of  a  body, 
b,  or  with  two  atoms  of  it,  or  with  three,  four,  etc.,  atoms 
of  it.  The  union  of  one  atom  of  a  with  one  of  b  produces 
one  compound,  the  union  of  one  atom  of  a  with  two  atoms 
of  b  produces  another  compound,  and  so  on." 

"We  have  no  means  of  demonstrating  the  number  of 
atoms  which  unite  together  in  this  manner  in  every  com- 
pound; we  must,  therefore,  have  recourse  to  conjecture. 
If  two  bodies  unite  only  in  one  proportion,  it  is  reasonable 
to  conclude  that  they  unite  atom  to  atom.  Hence  it  is 
most  likely  that  water  is  composed  of  one  atom  of  oxygen 
and  one  atom  of  hydrogen;  oxide  of  silver,  of  one  atom  of 
silver  and  one  atom  of  oxygen;  and  oxide  of  zinc,  of  one 
atom  of  zinc  and  one  atom  of  oxygen." 

"If  we  know  the  number  of  atoms  of  which  a  body  is 
combined,  and  the  proportion  of  the  constituents,  there  is 
no  difficulty  in  determining  the  proportional  weight  of  the 
atoms  of  which  it  is  composed.  Thus,  if  water  be  com- 
posed of  one  atom  of  oxygen  and  one  atom  of  hydrogen, 
and  if  the  weight  of  the  oxygen  in  water  is  to  that  of  the 
hydrogen  as  7^  to  1,  then  it  follows  that  the  weight  of  an 


GENERAL  CHEMISTRY.  79 

atom  of  oxygen  is  to  that  of  an  atom  of  hydrogen  as  7^ 
to  1." 

The  above  quotations  express  clearly  Dalton's  notion 
of  the  combination  of  atoms  to  form  larger  groups,  which 
we  now  call  molecules.  Because  of  lack  of  sufficient  ex- 
perimental data  he  was  led  to  assume  that  compounds 
formed  by  the  union  of  atoms  are  in  many  cases  simpler 
than  we  now'have  reason  to  consider  them,  but  this  is  a  detail 
which  does  not  detract  from  the  theory.  Dalton  recog- 
nized that  the  weights  of  these  minute  atoms  must  be 
exceedingly  small  and  beyond  the  reach  of  practical  deter- 
mination. He  therefore  proposed  a  new  system  of  weights, 
the  weight  of  the  atom  of  hydrogen  being  taken  as  the 
standard  and  called  unity.  The  weights  belonging  to  this 
system  are  commonly  called  the  atomic  weights,  and  one 
method  of  arriving  at  their  value  is  suggested  in  the  above 
quotation;  other  methods  will  be  pointed  out  later. 

Since  the  time  of  Dalton  many  new  elements  have  been 
discovered  and  exact  analytical  methods  have  been  per- 
fected by  which  the  weights  of  their  ultimate  atoms  on  the 
hydrogen  scale  may  be  readily  found.  For  certain  prac- 
tical reasons  Berzelius,  a  contemporary  of  Dalton,  sug- 
gested the  atom  of  oxygen  as  the  standard  and  proposed 
to  call  its  weight  arbitrarily  100.  This  suggestion  did  not 
meet  with  general  favor.  To-day,  however,  many  chem- 
ists agree  with  Berzelius  in  his  reasons  for  preferring  the 
atom  of  oxygen,  rather  than  the  atom  of  hydrogen,  as  the 
standard,  but  place  its  weight  at  16,  which  is  very  nearly 
its  true  weight  on  the  hydrogen  scale.  The  table  given 
below  contains  a  list  of  the  elements  with  their  atomic 
weights  on  the  hydrogen  scale  and  on  the  oxygen  scale. 
This  table  has  been  calculated  by  Prof.  F.  W.  Clarke,  of  the 
U.  S.  Geological  Survey,  and  embraces  the  results  of  the 
latest  and  most  accurate  determinations.  In  the  fifth  col- 
umn of  the  table  are  given  some  approximate  values  ob- 
tained by  rounding  off  the  numbers  of  the  fourth  column. 
These  approximate  values  are  convenient  and  sufficiently 
accurate  for  the  calculation  of  problems  to  be  given  later, 
and  also  for  the  illustrations  which  follow.  The  first  col- 
umn contains  the  name  of  the  element  and  the  second  col- 
umn the  symbol  by  which  it  is  represented. 


80 


GENERAL  CHEMISTRY. 
Table  of  Atomic  Weights. 


Name. 

Symbol. 

H=l 

O=16 

Approx. 

Aluminum 

Al 

26  91 

27  11 

27  1 

Antimony  

Sb 

119.52 

120.43 

120  4 

Argon    ....... 

Ar 

40  0 

Arsenic      ...          

As 

74  44 

75  01 

75  0 

Barium           

Ba 

136  39 

137.43 

137  4 

Beryllium  

Be 

9  01 

9  08 

Bismuth  ,  

Bi 

206.54 

208.11 

208  1 

Boron  

B 

10.86 

10.95 

11.0 

Br 

79.34 

79.95 

80  0 

Cd 

111.10 

111.95 

111.9 

CsGsium 

Cs 

131  89 

132  89 

Calcium  

Ca 

39.76 

40.07 

40.1 

Carbon                      

c 

11  92 

12  01 

12  0 

Cerium  .                           ... 

Ce 

139.1 

140  2 

Chlorine  

Cl 

35  18 

35  45 

35  5 

Chromium  .  .        .... 

Cr 

51.74 

52  14 

52  1 

Cobalt  

Co 

58.49 

58  93 

58  9 

Columbium  

Cb 

93.02 

93.73 

CoDDer 

Cu 

63  12 

63  60 

63  6 

Erbium  *i  

Er 

165.06 

166.32 

Fluorine 

F 

18  91 

19  06 

19  0 

Gadolinium 

Gd 

155  57 

156  76 

Gallium  

Ga 

69.38 

69  91 

Germanium  

Ge 

71.93 

72.48 

Gold  

Au 

195.74 

197.24 

197  2 

Helium  

He 

4  0 

Hydrogen  

H 

1.00 

1.008 

1.0 

Indium 

In 

112  99 

113  85 

Iodine  

I 

125.89 

126.85 

126.9 

Indium     

Ir 

191.66 

193  12 

Iron       .            

Fe 

55  60 

56  02 

56  0 

Lanthanum  

La 

137.59 

138  64 

Lead    

Pb 

205.36 

206  92 

206  9 

Lithium  

Li 

6.97 

7.03 

7  0 

Magnesium           . 

Me 

24  10 

24  28 

24  3 

Manganese                         

Mn 

54  57 

54  99 

55  0 

Mercury                 

Hg 

198.49 

200  00 

200  0 

Molybdenum  

Mo 

95.26 

95.98 

96  0 

Neodymium  

Nd 

139.70 

140.80 

Nickel  

Ni 

58.24 

58.69 

58.7 

Nitrogen               ... 

N 

13  93 

14  04 

14  0 

Osmium               .        .       

Os 

189.55 

190  99 

GENERAL  CHEMISTRY. 
Table  of  Atomic  Weights.— Continued. 


81 


Name. 

Symbol. 

H=rl 

O=16 

Approx. 

O 

15.88 

16.00 

16.0 

Pd 

105.56 

106.36 

Phosphorus    

P 

30.79 

31.02 

31.0 

Platinum  

Pt 

193.41 

194.89 

194.9 

K 

38.82 

39.11 

39.1 

Praseodymium                   .  ,  .. 

Pr 

142.5 

143  6 

Rhodium                                      .  . 

Rh 

102  23 

103  01 

Rubidium 

Rb 

84  78 

85  43 

Ruthenium                              

Ru 

100  91 

101  68 

Samarium                     

Sm 

149  13 

150  26 

Scandium                     .            ... 

Sc 

43  78 

44.12 

Selenium                          ... 

Se 

78.42 

79  02 

79 

Silicon                         

Si 

28  18 

28  4 

28  4 

Silver 

Ag 

107  11 

107.92 

107  9 

Sodium                                 . 

Na 

22  88 

23.05 

23  0 

Strontium                                      .  . 

Sr 

86  95 

87  61 

87.6 

Sulphur            o    

s 

31.83 

32  07 

32  1 

Tantalum                       

Ta 

181  45 

182.84 

Tellurium  

Te 

126  52 

127.49 

127  5 

Terbium       .    .                

Tb 

158  8 

160.0 

Thallium      ... 

Tl 

202.60 

204  15 

Thorium   

Th 

230  87 

282  63 

Thulium   

Tm 

169  4 

170.7 

Tin             

Sn 

118.15 

119  05 

119 

Titanium    

Ti 

47.79 

48.15 

Tungsten    

W 

183.43 

184.83 

Uranium  

u 

237  .  77 

239.59 

239  .  6 

Vanadium.    

V 

50.99 

51.38 

Ytterbium  

Yb 

171.88 

173.19 

Yttrium 

Yt 

88  35 

89  02 

Zinc 

Zn 

64  91 

65  41 

65  4 

Zirconium  

Zr 

89.72 

90.40 

Use  of  Symbols.  We  have  reached  a  point  now  in 
our  work  where  very  great  help  is  derived  from  the  use  of 
symbols  representing  the  substances  dealt  with,  and  the 
student  is  advised  to  learn  those  of  the  important  elements 
in  the  table. 

The  first  use  of  a  symbol  is  as  an  abbreviation  of  the  name 
of  a  substance  dealt  with.  Thus,  we  use  H  as  standing  for 


82  GENERAL  CHEMISTRY, 

or  representing  hydrogen,  O  as  representing  oxygen,  S  as 
representing  sulphur,  and  so  on.  Employed  in  this  man- 
ner, we  use  the  symbol  merely  to  save  time  or  space  in 
writing.  But  there  is  a  second  and  much  more  important 
application  of  these  letters  as  representing  something.  H 
stands  for  the  smallest  weight  of  the  element,  hydrogen, 
which  can  exist  in  any  compound  or  take  part  in  any  reac- 
tion, O  for  the  smallest  weight  of  oxygen,  S  for  the  small- 
est weight  of  sulphur  combining  or  existing  in  the  same 
manner.  According  to  the  definition  given  above,  these 
symbols,  therefore,  represent  weights  of  the  several  sub- 
stances corresponding  to  the  atomic  weights.  In  all  of  our 
calculations  O  stands  for  16  parts  of  oxygen,  H  for  1  part 
of  hydrogen,  Cl  for  35.5  parts  of  chlorine,  Na  for  23  parts 
of  sodium,  and  so  on. 

Molecules  and  Molecular  Weight.  The  atoms  of 
the  elements  mentioned  above  combine  among  themselves 
to  form  groups  called  molecules.  Two  atoms  of  hydrogen 
unite  with  one  atom  of  oxygen  to  form  a  molecule  of 
water;  an  atom  of  oxygen  forms  with  an  atom  of  mercury 
a  molecule  of  mercuric  oxide;  two  atoms  of  carbon,  six 
atoms  of  hydrogen  and  one  atom  of  oxygen  in  combina- 
tion constitute  a  molecule  of  alcohol.  We  represent  mole- 
cules by  uniting  the  symbols  of  their  component  atoms. 
Thus,  for  the  above  illustrations:  H2O,  HgO,  C2H6O. 

We  call  this  combination  of  symbols  a  formula.  A 
symbol  is,  therefore,  arbitrarily  taken  to  represent  an  atom, 
while  a  formula  represents  a  molecule. 

We  may  now  apply  some  of  these  facts  in  explanation 
of  experiments  in  the  preceding  chapters.  In  our  experi- 
ment on  the  decomposition  of  mercuric  oxide,  with  libera- 
tion of  oxygen,  we  may  represent  what  takes  place  by  this 
equation: 

HgO    :=   Hg    +    O 
MorxideC  =Mercury-f  Oxygen. 

This  equation  tells  us  that  216  parts  of  the  compound, 
mercuric  oxide,  yield  when  heated  200  parts  of  the  ele- 


GENERAL  CHEMISTRY.  83 

ment,  mercury,  and  16  parts  of  the  element,  oxygen.  In 
the  decomposition  of  potassium  chlorate  we  have  a  more 
complex  case.  This  is  a  combination  of  potassium,  oxygen 
and  chlorine  from  which  the  oxygen  may  be  readily  sep- 
arated by  heat.  A  study  of  the  compound,  potassium 
chlorate,  shows  that  it  contains  its  elements  in  these  pro- 
portions by  weight:  oxygen,  48  parts;  chlorine,  35.5  parts ; 
potassium,  39.1  parts.  The  atomic  weights  or  combining 
weights  are  represented  here,  three  times  for  oxygen,  once 
for  chlorine  and  once  for  potassium.  We  therefore  write 
as  the  formula  of  our  compound,  KC1O3.  When  we  de- 
compose this  we  find  that  all  of  the  oxygen  is  given  off 
and  that  we  have  a  solid  substance  left  which  contains 
potassium  and  chlorine  in  the  proportions  in  which  they 
were  found  in  the  original  compound.  We  therefore  write 
this  equation,  as  expressing  the  results  of  our  experiment : 

KC1O3    =      KC1    -f    O3 
Potassium     _  Potassium     I  Qxvgon 
chlorate       ~  chloride 

The  student  must  early  recognize  this  fact,  that  a  chem- 
ical equation  is  always  written  to  show,  in  compact  form, 
what  experiment  proves  has  taken  place  or  must  take  place 
under  proper  conditions.  Chemical  symbols  cannot  be 
combined  at  random,  we  cannot  perform  operations  on 
them  as  we  do  with  algebraic  symbols,  but  when  we  write 
them  on  the  left  hand  side  of  our  equality  sign  we  simply 
name  the  substance  or  substances  on  which  some  experi- 
ment is  to  be  performed.  After  the  experiment  we  are  able 
to  complete  the  equation,  and  then  we  write  down  on  the 
other  side  of  the  equality  sign  what  has  taken  place. 

The  above  equation  shows  correctly,  only  tne  relations 
by  weight  between  the  substance  taken  and  the  products. 
Careful  experiments  have  made  it  plain  that  the  reaction 
really  takes  place  in  two  stages,  as  represented  by  the  fol- 
lowing equations  : 


2KClO3^KClO4-fKCl+O2. 


84  GENERAL  CHEMISTRY. 

Potassium  chlorate  yields  at  first  a  substance  known  as 
potassium  perchlorate  with  potassium  chloride  and  a 
relatively  small  amount  of  oxygen,  in  fact  just  one-third, 
by  weight,  of  that  in  the  original  compound.  At  a  high 
temperature  the  potassium  perchlorate,  represented  by  the 
formula  KC1O4,  breaks  up  into  more  potassium  chloride 
and  oxygen,  thus, 

KC104:=KC1+04. 
Combining  the  two  equations  we  can  therefore  write: 

2KClO3  =  KCl-fKCl+O2+O4, 
or  better, 

2KC1O3=2KC1+3O2. 

Why  we  write  3O2  instead  of  O6  is  a  question  which 
cannot  be  satisfactorily  answered  at  this  point,  but  will  be 
taken  up  later  in  consideration  of  other  experimental  re- 
sults; but  another  simple  matter  in  connection  with  the 
method  of  writing  equations  must  be  explained  here.  The 
student  observes  that  numerals  are  employed  in  two  posi- 
tions in  these  equations.  Large  figures  are  written  before 
the  formulas  of  compounds,  while  small  figures  in  several 
cases  seem  to  be  to  the  right  and  a  little  below  certain 
symbols.  This  is  a  purely  conventional  arrangement,  and 
the  meaning  conveyed  could  be  just  as  well  expressed  in 
some  other  manner.  It  has  been  agreed  by  chemists  to 
consider  the  large  numerals  as  multiplying  the  whole  com- 
pound which  follows,  while  the  small  figures  are  taken  as 
referring  only  to  the  symbol  of  the  element  immediately 
preceding.  Thus,  in  2KC1O3  we  have  twice  the  whole 
group,  while  the  small  3  indicates  that  we  have  in  each 
group  the  combining  weight  of  oxygen  taken  three  times. 
The  effect  of  the  large  numerals,  however,  is  not  carried 
beyond  a  sign  of  addition  or  subtraction.  In  the  equation, 


the  large  2  at  the  beginning  refers  only  to  the  KNO3,  and 
not  to  H2SO4. 


GENERAL  CHEMISTRY.  85 

We  may  now  express,  by  the  use  of  symbols,  the  reac- 
tions which  took  place  in  some  of  the  experiments  on  oxy- 
gen. We  have 

C    +    02   =  C02. 

Carbon  +  Oxygen  =  Carbon 

S    +    O2   =  SO2. 

Sulphur  +  Oxygen^ 

P4 

Phosphorus 


=    2P206. 


In  the  first  case  we  have  the  union  of  12  parts  of  car- 
bon with  32  parts  of  oxygen.  In  the  second  case  32.1  parts 
of  sulphur  combine  with  32  parts  of  oxygen,  while  in  the 
third  case  124  (4X31)  parts  of  phosphorus  combine  with 
80  (5  X  16)  parts  of  oxygen. 

The  combining  weight  of  zinc  has  been  found  to  be 
65.4,  and  we  find  that  65.4  parts  of  zinc  act  on  98.1  parts 
of  sulphuric  acid,  liberating  2  parts  of  hydrogen.  We  there- 
fore write  the  equation, 


Zn  -fH2S04  =  ZnS04  +  H2. 

Zinc  +>SfC  =su^hCate     +  Hydrogen. 


This  equation,  like  all  the  others,  is  intended  to  express 
the  results  of  experiments.  We  find  that  sulphuric  acid  is 
composed  of  hydrogen,  oxygen  and  sulphur  in  certain  pro- 
portions shown  by  the  formula  H2SO4,  and  that  at  the  end 
of  the  experiment  we  have  in  solution  a  compound  which 
contains,  in  place  of  hydrogen,  zinc  combined  with  oxygen 
and  sulphur,  in  the  proportions  shown  by  the  formula, 
ZnSO4,  and  which  is  known  as  zinc  sulphate.  By  the  sym- 
bol H  we  represent  1  combining  weight  of  hydrogen,  and 
by  H2  we  represent  two  such  weights.  By  our  equation, 
therefore,  we  express  this  fact,  that  in  the  solution  of  65.4 
parts  of  zinc  we  liberate  2  parts  by  weight  of  hydrogen. 
To  generate  2  grams  of  hydrogen  gas  by  this  method  we 
must  dissolve  65.4  grams  of  the  metal  zinc. 


86  GENERAL  CHEMISTRY. 

It  has  been  already  mentioned  that  we  may  use  iron  in- 
stead of  zinc  for  the  generation  of  the  hydrogen  gas.  In 
this  case  we  find  that  56  parts  of  iron  produce  the  same 
amount  of  hydrogen  that  we  obtain  from  the  65.4  parts  of 
zinc.  The  56  represents,  in  fact,  the  combining  weight  of 
iron,  and  we  may  write  as  expressing  the  last  reaction, 


Fe  +H2SO4  =  FeSO4-f     H2 

Tr/,»,   I   Sulphuric          Ferrous     (    „   j 
Iron+      acid         =  sulphate  +  Hydrogen. 

56  +   98.1      =    152.1  -f-       2. 

The  sum  of  the  weights  on  the  left  hand  side  of  the 
equation  is  equal  to  the  sum  on  the  right.  There  is  neither 
a  gain  nor  loss  of  matter,  but  merely  a  rearrangement  of 
elements  in  the  compounds.  It  appears  from  these  illus- 
trations that  zinc  and  iron  have  the  power  of  displacing  the 
hydrogen  in  the  acid  used.  Had  we  used  hydrochloric  acid 
instead  of  sulphuric  exactly  the  same  behavior  would  have 
been  observed.  Hydrogen  would  have  been  displaced  in 
quite  the  same  manner,  and  2  parts  by  weight  for  65.4 
parts  of  zinc,  or  56  parts  of  iron  dissolved.  In  general,  it 
may  be  said  that  zinc  and  iron  displace  hydrogen  in  many 
bodies  called  acids,  and  always  in  the  proportions  given. 

We  come  now  to  a  consideration  of  the  reactions  by 
which  chlorine  and  hydrochloric  acid  were  produced,  and 
here  again  we  deal  with  the  results  of  exact  experiments. 
It  has  been  found  that  sulphuric  acid  is  able  to  decompose 
sodium  chloride,  or  common  salt,  in  a  manner  illustrated 
by  the  following  equation: 


Na2SO4-h    2HC1 

Sulphuric  ,     Sodium    ___      Sodium      i    Hydrochloric 
acid      ~t~  chloride  ~  "     sulphate    "1  acid 

98.1  +     117     =    142.1     -f        73. 

Experiment  actually  shows  us  that  for  an  amount  of 
sulphuric  acid  represented  by  the  sum  of  the  combining 
weights  of  its  elements,  that  is  for  98.1  parts  by  weight, 


GENERAL  CHEMISTRY.  87 

we  require  117  parts  of  salt,  that  is,  twice  the  sum  of  the 
combining  weights  of  the  sodium  and  chlorine. 

When  we  employ  the  pure  materials  in  exactly  these 
proportions,  and  aid  their  action  on  each  other  by  heat,  we 
find  at  the  end  of  our  experiment  that  we  have  neither 
sulphuric  acid  nor  salt,  but  two  new  substances,  one  of 
which  is  the  hydrochloric  acid,  which  we  collect  in  water, 
and  the  other  a  white  solid  substance  which  remains  in  the 
decomposing  vessel,  and  which  we  call  sodium  sulphate. 
There  are  no  other  products  in  the  reaction.  We  call  this 
a  reaction  of  double  decomposition,  inasmuch  as  we  start  with 
two  compound  bodies  which  react  on  each  other  to  form 
two  new  compound  bodies.  We  can  illustrate  this  double 
decomposition  by  a  diagram,  as  follows: 


C-        -D  C^  "ND 

Before  the  reaction,  one  compound  bod}7  is  made  up  of 
the  parts  A  and  B,  and  the  other  compound  body  of  the  parts 
C  and  D.  But  after  the  decomposition  we  have  a  new 
compound  body,  with  A  and  D  as  its  parts,  and  another 
with  C  and  B.  The  reaction  between  salt  and  sulphuric 
acid  is  a  typical  one  of  double  decomposition  and  well 
illustrates  many  which  are  to  follow. 

From  our  hydrochloric  acid,  as  made  above,  we  sepa- 
rate the  chlorine  by  another  process  described.  This  is 
somewhat  more  complex,  but  its  exact  nature  may  be  read- 
ily illustrated  as  follows: 


MnO2+     4HC1*    s±  MnCl2  +2H8O+     C12 

Manganese  _1_  Hydrochloric  —  Manganese_|_    Wiff>r    4-rhlnrinp 
dioxide  acid  ~    chloride      ' 

87      -f       146       =      126     -f     36     -f      71. 

The  manganese  dioxide  in  the  above  is  made  up  of  the 
combining  weight  of  manganese  plus  twice  that  of  oxygen. 
Now  experiment  shows  that  to  decompose  this  completely 


88  GENERAL  CHEMISTRY. 

we  need  of  hydrochloric  acid  four  times  the  sum  of  the 
combining  weights  of  hydrogen  and  chlorine.  Less  would 
not  be  sufficient  to  complete  the  decomposition  of  the 
manganese  dioxide.  The  experiment  may  be  performed 
in  such  a  manner  as  to  show  that  water  is  liberated,  and 
exactly  how  much.  For  the  amount  of  manganese  dioxide 
assumed  to  be  taken  we  find  of  water  just  twice  the  sum 
of  the  combining  weights  of  hydrogen  and  oxygen,  the 
weights  taken  or  obtained  being  expressed  in  any  con- 
venient standard,  as  in  grams.  The  chlorine  liberated  is 
just  half  of  that  contained  in  the  original  hydrochloric 
acid  used,  which  fact  is  expressed  also  in  the  equation. 

When  chlorine  is  liberated  by  the  action  of  salt,  sul- 
phuric acid  and  manganese  dioxide  on  each  other,  the 
whole  of  that  element  in  the  salt  may  be  obtained.  The 
reaction  takes  place  in  two  stages  possibly,  the  first  involv- 
ing the  formation  of  hydrochloric  acid  and  the  second  its 
decomposition.  The  following  equation  shows  the  quanti- 
tative relations  existing  between  the  compounds  taken 
and  obtained. 


+  2NaCl+2H2S04  =  MnS04+Na2S04  +  2 
87     +     11?    +    196.2     =    151.1    +    142-1     +  36       +71. 

From  the  above  it  appears  that  the  manganese  dioxide 
used,  and  chlorine  obtained  stand  to  each  other  in  the  re- 
lation of  87  to  71.  As  both  weights  are  referred  to  the 
same  basis  or  standard,  the  proportion  must  hold  good  if 
we  refer  them  to  a  new  standard.  If  we  take  the  gram  as 
our  unit  it  is  true  that  87  grams  of  manganese  dioxide 
must  be  used  for  the  liberation  of  71  grams  of  chlorine, 
and  if  we  weigh  in  pounds  ortons.the  same  ratio  must  still 
exist.  It  is  easy,  therefore,  to  tell  how  much  manganese 
dioxide  must  be  used  to  liberate  any  given  quantity,  as 
100  grams  of  chlorine.  It  is  evident  that  the  correct 
answer  must  be  given  by  the  following  proportion  : 

87  :  71  ::  x  :  100. 
x  =  122.5  grams. 


GENERAL  CHEMISTRY.  89 

To  find  the  amount  of  salt  required  we  make  another 
proportion  : 

117  :  71  ::  x  :  100. 

x=164.8  grams. 

A  third  proportion  shows  the  amount  of  sulphuric  acid 
necessary  for  the  formation  of  the  100  grams  of  chlorine  : 

196.2  :  71  ::  x  :  100. 
x  =  276.3  grams. 

Aided  by  the  above  explanations  the  student  should 
now  be  able  to  understand  what  follows.  From  this  point 
on  all  important  reactions  will  be  represented  by  equations, 
and  these  should  be  thoroughly  studied.  The  student 
should  keep  in  mind,  however,  that  equations  are  not 
drawn  from  the  imagination,  but  represent,  properly,  the 
results  of  experiments.  He  should  practice  writing  them 
as  an  aid  to  memorizing  important  reactions,  and 
especially  because  of  their  value  in  the  solution  of  even 
the  simplest  chemical  problems,  as  illustrated  by  the 
examples  given  above. 

In  the  present  chapter  no  attempt  will  be  made  to 
explain  methods  by  which  the  atomic  weights  are  found. 
In  one  to  follow,  however,  after  the  student  has  become 
more  familiar  with  chemical  facts,  something  on  this  topic 
will  be  given. 


CHAPTER   IV. 


COMPOUNDS  OF    CHLORINE  WITH    OXYGEN.-  BRO- 
MINE, IODINE,  FLUORINE  AND  THEIR  COMPOUNDS. 

TN   the  last  chapter  the  element,  chlorine,  and   its  com- 
pound   with    hydrogen    have  been  described.     In  this 
chapter    a    few    other   important    combinations    must    be 
referred  to. 

OXIDES  AND  ACIDS  OF  CHLORINE. 

Three  compounds  of  chlorine  with  oxygen  are  known, 
but  they  cannot  be  formed  by  direct  union. 

Chlorine  Monoxide  and  Hypochlorites.  The  first 
one  is  called  chlorine  monoxide,  and  is  represented  by 
the  formula  C12O.  It  is  a  yellowish  brown  gas  with  an 
odor  suggesting  chlorine,  and  may  be  made  by  passing 
dry  chlorine  over  the  red  oxide  of  mercury,  freshly  pre- 
cipitated and  dried,  and  contained  in  a  glass  tube.  This 
equation  expresses  the  combination  : 


The  gas  can  be  easily  condensed  to  a  liquid,  but  this 
is  not  stable.  It  is  very  soluble  in  water,  forming  a  new 
acid,  called  hypochlorous  acid  : 

C12O  +  H2O  =  2HOC1. 
The  acid  solution  is  not  stable;  if  the  gas  is  led  into  an 


GENERAL  CHEMISTRY.  91 

alkali  solution,  however,  an  important  body  called  a  hypo- 
chlorite  is  formed: 

C13O+2KOH  =  2KOC1-|-H2O. 

Some  of  the  hypochlorites  are  well  known  and  useful 
substances.  Calcium  hypochlorite  is  the  active  constitu- 
ent of  bleaching  powder.  Sodium  and  potassium  hypochlo- 
rites are  used  in  the  laboratory  and  in  medicine.  These 
hypochlorites  are  easily  decomposed  by  hydrochloric  or 
sulphuric  acid  with  liberation  of  chlorine.  Practical  appli- 
cation is  made  of  this  in  bleaching  by  bleaching  powder, 
which  is  illustrated  by  the  following  experiment: 

Ex.  72.  Pour  some  dilute  sulphuric  acid  over  a  few  grams  of 
bleaching  powder  in  the  bottom  of  a  large  beaker  which  then  cover  with 
a  piece  of  glass  or  a  card.  Observe  that  greenish  yellow  fumes  soon 
collect  in  the  beaker.  Moisten  now  a  piece  of  bright  calico,  as  already 
described,  and  hang  it  in  the  beaker  of  gas.  The  calico  will  fade  as 
before. 

The  manufacture  of  bleaching  powder  will  be  referred 
to  later.  It  is  made  by  passing  chlorine  gas  over  slaked 
lime,  and  is  essentially  a  mixture  of  calcium  hypochlorite, 
CaO2Cl2,  and  calcium  chloride,  CaCl2,  in  about  equal 
proportions. 

Chlorine  Dioxide.  This  is  a  heavy  dark  yellow  gas 
usually  made  by  the  decomposition  of  potassium  chlorate 
by  sulphuric  acid.  Chloric  acid,  HC1O3,  is  formed  first 
and  this  decomposes  on  slight  warming. 


The  gas  can  be  condensed  to  a  liquid  at  a  low  tempera- 
ture. It  is  not  stable,  often  decomposing  with  explosive 
violence.  It  dissolves  rather  readily  in  water  but  does  not 
form  a  new  acid. 

Chlorine  Trioxide.  This  is  a  greenish  yellow  gas 
having  the  composition,  C12O3,  and  is  made  by  several  proc- 
esses depending  on  the  reduction  of  chloric  acid,  HC1O3. 


92  GENERAL  CHEMISTRY. 

It  is  decomposed  by  warm  water  forming  a  mixture  of 
hydrochloric  and  chloric  acids,  and  with  cold  water 
yields  chlorous  acid,  as  below. 

Chlorous  Acid,  HC1O2,  is  not  known  in  the  pure 
state,  but  certain  salts,  called  chlorites,  are  known  and  these 
correspond  to  the  acid. 

Solutions  of  chlorous  acid  result  when  the  trioxide  is 
dissolved  in  water: 

C12O3+H2O  =  2HC102. 

Chloric  Acid.  This  is  the  best  known  of  the  oxygen 
acids  of  chlorine  and  may  be  prepared  by  decomposing 
barium  chlorate  by  means  of  sulphuric  acid  : 

Ba(ClO3)2  +  H2SO4=:2HClO3-hBaSO4. 

As  barium  sulphate,  BaSO4,  is  a  very  insoluble  precip- 
itate it  is  easy  to  obtain  a  pure  solution  of  the  chloric 
acid  by  pouring  off  the  supernatant  liquid.  The  con- 
centrated acid  is  a  strong  oxidizing  agent,  has  a  pungent 
odor  and  decomposes  readily  when  heated,  yielding 
a  new  acid  known  as  perchloric  acid,  HC1O4,  along 
with  oxygen  and  chlorine.  Corresponding  to  chloric 
acid,  we  have  the  well-known  salts  called  chlorates,  of 
which  potassium  chlorate,  KC1O3,  is  the  best  illustration. 
The  chlorates  are  all  soluble  in  water  and  decompose 
when  heated,  yielding  oxygen.  The  decomposition  takes 
place  in  two  stages,  however;  in  the  first  perchlorate  is 
formed: 

2KC1O3=KC104 


In  the  second  stage  of  the  eaction  the  perchlorate  is 
decomposed,  yielding  more  oxygen  and  chloride. 

Perchloric  Acid.  As  potassium  perchlorate  is  but 
slightly  soluble  in  water,  advantage  is  taken  of  the  above 
reaction  in  preparing  perchloric  acid,  HC1O4.  When  the 


GENERAL  CHEMISTRY.  93 

chlorate  is  heated  until  the  evolution  of  oxygen  begins  the 
first  stage  of  the  reaction  may  be  considered  as  completed, 
practically.  If  the  mass  is  now  cooled,  powdered  and 
extracted  with  water  the  perchlorate  is  left  while  the 
chloride  goes  into  solution.  This  perchlorate  distilled 
with  strong  sulphuric  acid  yields  perchloric  acid,  which  is  a 
heavy,  volatile  liquid  having  a  great  affinity  for  water.  It 
is  a  powerful  oxidizing  agent  and  decomposes  immediately 
when  brought  in  contact  with  most  organic  substances. 
The  perchlorates  are  all  soluble  in  water  and  they  differ 
from  the  chlorates  in  not  being  decomposed  by  hydro- 
chloric acid. 

It  appears  from  the  foregoing  that  we  have  four  chlorine 
acids  containing  oxygen.  The  names  and  formulas  of  these 
are  : 

Hypochlorous  acid,  HC1O, 
Chlorous  acid,  HC1O2, 
Chloric  acid,  HC1O3, 
Perchloric  acid,  HC1O4. 

It  will  be  observed  that  the  names  differ  through  certain 
prefixes  and  terminations  and  it  will  be  seen  later  that  the 
same  are  used  in  the  designations  of  all  other  acids. 

OTHER  CHLORINE  COMPOUNDS. 

Chlorine  combines  indirectly  with  nitrogen  to  form  a 
very  explosive  substance  known  as  nitrogen  chloride.  It 
forms  a  number  of  important  combinations  with  carbon  and 
with  carbon  and  hydrogen,  to  be  mentioned  later.  A  very 
important  compound  with  oxygen  and  nitrogen  will  be 
described  in  the  next  chapter. 

BROMINE. 

Occurrence.  Bromine  is  an  important  element  which 
never  occurs  free  in  nature.  It  is  found  in  several  bro- 
mides in  spring  waters,  and  to  a  slight  extent  in  sea  water. 


94  GENERAL  CHEMISTRY. 

History.  Bromine  was  discovered  in  1826  by  Balard 
in  the  mother  liquor  left  after  crystallization  of  salt  from 
evaporated  sea  water.  The  discoverer  was  able  to  show 
the  important  analogies  existing  between  this  element  and 
chlorine  and  iodine. 

Preparation.  Much  of  our  bromine  is  obtained  from 
the  residues  left  on  crystallizing  salt  from  concentrated 
brine  of  certain  salt  springs.  The  bromine  is  left  in  these 
mother  liquors  in  the  form  of  bromides,  which  are  more 
soluble  than  the  common  salt,  and  may  be  liberated  by 
several  reactions,  of  which  two  illustrations  will  be  given. 
Large  quantities  of  bromine  are  produced  at  the  Michigan 
salt  wells  and  also  from  the  salt  deposits  of  Stassfurt, 
Germany. 

Ex.  73.  Dissolve  a  small  crystal  of  sodium  or  potassium  bromide 
in  water  in  a  test-tube,  and  add  gradually,  a  drop  at  a  time,  some 
chlorine  water.  Use  for  this  purpose  the  chlorine  water  saved  from  a 
former  experiment.  When  the  first  drop  of  chlorine  water  mixes  with 
the  solution  of  bromide  it  produces  a  reddish-yellow  color  which 
deepens  to  red  as  more  of  the  reagent  is  added.  If  the  bromide  solution 
is  weak  and  the  chlorine  water  strong,  the  red  color  will  finally  disap- 
pear by  continued  addition  of  the  latter.  The  chlorine  water  decom- 
poses the  bromide,  liberating  bromine. 

In  the  above  experiment  we  have  illustrations  of  sev- 
eral important  points.  First,  of  the  liberation  of  free  bro- 
mine. Potassium  bromide  is  a  combination  of  bromine 
with  potassium,  which  we  represent  by  the  formula  KBr. 
Sodium  bromide  is  represented  by  NaBr.  Assuming  that 
we  are  dealing  with  the  former  we  express  the  whole 
reaction  by  this  equation: 

KBr  -f     Cl    =   KC1    +   Br 
119.1  +  35-5  =  V4.60  -f-    80. 

The  equation  shows  just  what  an  exact  quantitative  ex- 
periment would  have  revealed  to  us,  viz.,  that  35.5  parts 
by  weight  of  chlorine  are  required  to  completely  decompose 
119.1  parts  of  potassium  bromide  with  liberation  of  80 
parts  of  bromine.  The  weight  of  chlorine  taken  and  that 
of  bromine  obtained  are  chemically  equivalent,  but,  as  the 


GENERAL  CHEMISTRY.  95 

result  shows,  the  chlorine  is  able  to  displace  the  bromine. 
We  are  no  more  able  to  give  an  exact  reason  for  this  dis- 
placement than  we  are  to  assign  a  reason  for  other  chem- 
ical decompositions  already  illustrated.  But  we  are  accus- 
tomed to  say  that  the  chlorine  has  a  greater  affinity  for 
the  potassium  than  the  bromine  has  and  is  therefore  able 
to  drive  it  out  from  its  combination.  Of  the  real  nature 
of  this  chemical  affinity  we  know  but  little. 

The  above  experiment  illustrates  the  marked  activity 
of  chlorine  in  another  manner.  It  was  shown  that  a 
great  excess  of  the  chlorine  water  discharged  the  color  of 
the  free  bromine.  This  loss  of  color  is  due  to  two  causes. 
First,  to  the  combination  of  the  excess  of  chlorine  with  the 
bromine  liberated,  forming  bromine  chloride,  and  second, 
to  the  oxidation  of  some  of  the  bromine  to  bromic  acid,  in 
presence  of  water,  which  is  illustrated  by  this  equation: 

Br+3H2O  +  5Cl  =  5  HCl+HBrO3. 

Hydrochloric  and  bromic  acids  result. 

Bromine  can  be  readily  liberated  from  bromides  by  a 
reaction  analogous  to  that  employed  for  the  preparation  of 
chlorine  from  chlorides,  that  is,  by  the  use  of  sulphuric 
acid  and  manganese  dioxide.  The  following  experiment 
will  illustrate  this: 

Ex.  74.  In  a  300  Cc.  flask  mix  about  2  Gm.  of  powdered  potassium 
bromide  with  4  or  5  Gm.  of  commercial  powdered  manganese  dioxide. 
Add  a  little  water  and  shake  until  the  mixture  becomes  uniformly  dis- 
tributed. Then  add  50  Cc.  of  dilute  sulphuric  acid  and  close  the  flask 
with  a  stopper  through  which  passes  a  long  delivery  tube  bent  down  to 
dip  into  a  small  flask  or  beaker  of  cold  water.  The  flask  with  the  above 
described  mixture  must  stand  on  a  sand-bath  or  wire  gauze,  which  is 
then  heated  by  a  lamp.  Red  vapors  are  generated  in  the  flask  which 
distill  over  and  dissolve  in  the  water  in  the  small  receiving  flask  or 
beaker.  Continue  the  application  of  heat  as  long  as  these  red  vapors 
are  evolved.  Then  remove  the  receiving  flask  and  withdraw  the  lamp 
from  the  other  flask. 

The  reaction  which  takes  place  here  is  illustrated  by 
this  equation: 

2KBr+MnO24-2H2SO4=: 

Br24-MnSO4-f-K2SO4+2H2O. 


96  GENERAL  CHEMISTRY. 

This  is  seen  to  be  similar  to  the  chlorine  reaction.  A 
very  large  excess  of  manganese  dioxide  is  taken  in  the  ex- 
periment, in  order  to  secure  the  complete  decomposition 
of  the  bromide  without  liberation  of  hj'drobromic  acid. 

On  the  large  scale  this  process  is  applied  to  the  manu- 
facture of  bromine  from  the  mother  liquors  of  salt  works. 
As  these  liquors  contain  much  chloride,  chlorine  is  first 
liberated  and  this  serves  to  free  the  bromine.  Some 
bromine  chloride  is  always  produced  in  the  operation,  but 
this  is  more  volatile  than  the  bromine  and  can  be  sepa- 
rated by  distillation. 

In  the  experiment  just  described,  bromine  was  collected 
in  water.  It  is  somewhat  soluble  in  water,  as  shown  by 
the  fact  that  at  first  all  that  distilled  over  went  into  solu- 
tion. Before  the  end  of  the  experiment,  however,  unless 
too  much  water  was  taken  a  part  of  the  bromine  settles 
out  as  a  dark  red  drop.  Use  this  aqueous  solution  of  bro- 
mine for  tests  as  follows: 

Ex.  75.  Bromine  bleaches  as  does  chlorine,  but  with  less  activity. 
Test  this  by  use  of  colored  calico,  and  also  with  solutions  of  organic  col- 
oring matters,  litmus  and  cochineal,  for  instance. 

Bromine  is  readily  soluble  in  chloroform,  carbon  disul- 
phide,  ether  and  other  liquids,  and  may  be  withdrawn 
from  aqueous  solution  by  them. 

Ex.  76.  Pour  some  of  the  bromine  water,  made  above,  into  a 
test-tube,  and  add  about  one-tenth  its  volume  of  carbon  disulphide. 
Close  the  tube  with  a  cork  and  shake  thoroughly.  On  standing,  the 
disulphide  speedily  collects  at  the  bottom  of  the  tube,  and  it  will  be 
seen  that  it  is  highly  colored  by  the  absorbed  bromine,  while  the  water 
above  is  much  lighter  colored  than  before,  or  it  may  be  even  colorless. 
This  beautiful  reaction  is  employed  in  the  detection  of  small  traces  of 
bromides  in  spring  water,  the  bromine  being  first  liberated  by  means  of 
a  small  amount  of  chlorine. 

Physical  Properties.  Bromine  boils  at  63°,  and 
freezes  about  — 7°.  At  0°  it  has  a  specific  gravity  of  3. 18. 
At  15°  it  dissolves  in  33  parts  of  water. 

Uses.  Bromine  is  employed  in  the  preparation  of 
bromides,  several  of  which  are  used  in  medicine.  It  is 


GENERAL  CHEMISTRY.  9? 

used  also  in  making  certain  reagents  employed  in  labora- 
tories and  in  making  a  number  of  valuable  organic  prepa- 
rations. 

BROMINE  AND    HYDROGEN. 

Under  certain  conditions  these  two  elements  may  be 
directly  united,  but  not  as  readily  as  is  the  case  with  chlo- 
rine and  hydrogen.  Hydrobromic  acid,  HBr,  results. 
This  acid  cannot  be  made  in  pure  condition  by  the  reaction 
employed  in  the  manufacture  of  hydrochloric  acid,  that  is, 
by  the  decomposition  of  a  bromide  by  means  of  strong  sul- 
phuric acid,  according  to  the  following  equation  : 


The  hydrobromic  acid  as  liberated  is  partially  decom- 
posed by  the  excess  of  strong  sulphuric  acid,  free  bromine 
and  sulphurous  oxide  being  formed.  This  can  be  illus- 
trated as  follows  : 

Ex.  77.  Take  some  small  crystals  of  potassium  bromide  in  a  test- 
tube,  and  pour  over  them  a  little  strong  sulphuric  acid.  An  escape  of 
gas  is  seen  to  follow,  which  has  a  yellowish  color,  due  to  fi'ee  bromine, 
the  hydrobromic  acid  itself  being  a  colorless  gas. 

Pure  hydrobromic  acid  may  be  made,  however,  by  using 
phosphoric  acid  instead  of  sulphuric  acid,  and  by  a  reaction 
illustrated  by  this  equation  : 

PBr3+3H2O  =  3HBr-{-H3PO3. 

Phosphorous  bromide  is  decomposed  by  water,  yielding 
hydrobromic  acid  and  phosphorous  acid.  Instead  of  using 
pure  phosphorous  bromide  it  is  customary  to  add  bromine 
very  slowly  to  a  mixture  of  red  phosphorus  and  water  in  a 
suitable  apparatus  arranged  in  such  a  manner  that  the  gas 
as  it  escapes  may  be  absorbed  in  water. 

Properties.  In  most  of  its  important  properties  hy- 
drobromic acid  resembles  hydrochloric  acid.  It  is  a  gas, 
and  very  soluble  in  water.  One  cubic  centimeter  weighs 
.003645  Gm.  at  0°  and  normal  pressure.  It  is  readily  de- 


98  GENERAL  CHEMISTRY. 

composed  by  chlorine  and   it  unites  with  alkalies  forming 
bromides. 

Hydrobromic  acid  and  all  the  soluble  bromides  give  a 
precipitate  when  treated  with  a  solution  o'f  silver  nitrate  as 
shown  below  : 

Ex.  78.  Prepare  a  dilute  solution  of  potassium  bromide  by  dis- 
solving a  small  crystal  in  water.  Add  to  this  solution  a  few  cubic  cen- 
timeters of  a  dilute  solution  of  silver  nitrate.  A  yellowish  white,  curdy 
precipitate  forms  which  soon  settles  to  the  bottom  of  the  vessel  in  which 
it  was  produced.  This  precipitate  is  not  soluble  in  nitric  acid,  and  to 
a  limited  extent  only,  in  dilute  ammonia. 

The  compounds  of  bromine  with  hydrogen  and  oxygen 
are  not  of  sufficient  importance  to  be  taken  up  in  this 
place  in  detail.  No  oxides  are  known,  but  two  oxygen 
acids,  hypobromous  acid,  HBrO,  and  bromic  acid,HBrO3, 
are  known.  Some  of  the  hypobromites,  the  salts  formed 
from  HBrO,  are  used  as  reagents. 


IODINE. 

Iodine  is  a  very  important  element  resembling  chlorine 
and  bromine  in  certain  chemical  properties,  but  is  a  steel 
gray  solid  at  the  ordinary  temperature.  It  is  far  less 
abundant  than  either  chlorine  or  bromine  in  nature,  occur- 
ring in  some  springs,  but  mainly  in  sea  water,  from  which 
it  is  taken  up  by  certain  seaweeds.  The  iodine  of  com- 
merce is  largely  obtained  from  the  ash  produced  by  burn- 
ing these  weeds.  It  occurs  also  in  small  amount  in  com- 
pounds called  iodates,  which  occur  with  Chili  saltpeter. 

History.  Iodine  was  discovered  by  Courtois,  a  French 
chemist,  in  1812.  It  was  found  in  the  mother  liquors  left 
after  the  extraction  of  sodium  salts  from  kelp  or  the  ash 
of  seaweeds.  Until  recently  this  kelp,  or  varec,  was  the 
source  from  which  practically  all  iodine  was  obtained.  A 
larger  proportion  is  now  produced  from  the  mother  liquors 
occurring  in  the  refining  of  Chili  saltpeter.  As  produced 
from  seaweed  iodine  is  obtained  mainly  from  the  coasts  of 
Scotland  and  northern  France. 


GENERAL  CHEMISTRY.  99 

Preparation.  As  found  in  the  ash  from  seaweed,  the 
iodine  occurs  in  the  form  of  an  iodide  and  can  be  sep- 
arated just  as  bromine  is  from  bromides.  The  following 
experiments  will  illustrate  this  : 

Ex.  79.  Dissolve  a  small  crystal  of  potassium  iodide  in  about  five 
Cc.  of  water  in  a  test-tube.  Then  add  chlorine  water,  a  drop  at  a  time, 
which  produces  a  brown  color,  and  finally,  if  the  solution  is  not  too 
weak,  a  precipitate  of  free  iodine.  An  excess  of  chlorine  water  dis- 
charges the  color  as  in  the  corresponding  case  with  a  bromide,  and  for 
the  same  general  reasons. 

The  decomposition  is  illustrated  by  this  equation  : 


35.5  parts  of  chlorine  replace  126.9  parts  of  iodine. 

Iodine  is  liberated,  also,  by  the  reaction  with  manga- 
nese dioxide  and  sulphuric  acid,  which  can  be  easily  illus- 
trated by  a  simple  experiment  as  follows  : 

Ex.  80.  Mix  about  a  gram  of  powdered  potassium  iodide  with  two 
or  three  times  this  weight  of  powdered  manganese  dioxide  in  a  flask  of 
300  to  400  Cc.  capacity.  Pour  in  5  Cc.  of  dilute  sulphuric  acid  and 
heat  the  flask  on  a  sand-bath.  Decomposition  of  the  iodide  takes  place 
and  deep  violet  colored  vapors  fill  the  flask.  The  vapors  condense,  in 
.part,  on  the  upper  and  cooler  portions  of  the  flask. 

The  decomposition  is  illustrated  by  the  equation, 


2. 


K2SO4-fMnSO4-f2H2O-fI 


On  the  large  scale  the  reaction  is  so  carried  out  that 
the  iodine  distills  over  from  the  decomposing  retorts  of 
earthenware  or  firebrick,  and  condenses  in  cold  receivers. 

Properties.  Commercial  iodine  occurs  as  a  steel  gray 
crystalline  solid.  In  its  power  of  combination  with  metals 
iodine  is  less  active  than  chlorine  or  bromine.  The 
important  properties  of  the  substance  may  be  shown  by 
simple  experiments. 


100  GENERAL  CHEMISTRY. 

Ex.  81.  Heat  a  small  crystal  of  iodine  in  a  test-tube.  The  iodine 
vaporizes  quickly,  so  that  the  whole  tube  may  be  filled  with  the  violet 
colored  vapors.  That  these  vapors  are  heavier  than  air  may  be  shown 
by  holding  the  tube  in  front  of  a  sheet  of  white  paper  as  a  background, 
and  then  turning  the  tube  so  that  the  vapor  may  flow  down  and  show 
against  the  paper.  When  the  tube  cools,  add  a  little  distilled  water  and 
shake  thoroughly.  Iodine  is  slightly  soluble  in  water,  which  is  shown 
by  the  yellow  color  imparted  in  this  test.  To  the  aqueous  solution  add 
a  few  drops  of  cold  starch  paste.  A  beautiful  blue  results,  due  to  the 
combination  of  the  starch  with  iodine. 

Ex.  82.  Iodine  is  much  more  soluble  in  alcohol  than  in  water. 
Powder  a  small  crystal  of  iodine  and  transfer  to  a  test-tube.  Then  add 
about  2  Cc.  of  alcohol  and  shake  thoroughly.  The  iodine  dissolves, 
producing  a  brown  solution  known  as  the  tincture  of  iodine.  A  drop  of 
this  tincture  added  to  a  beaker  of  water  containing  a  little  starch  paste 
produces  a  blue  color. 

Ex.  83.  Iodine  dissolves  very  readily  in  an  aqueous  solution  of 
potassium  iodide,  which  can  readily  be  shown  by  adding  a  small  amount 
of  powdered  iodine  to  potassium  iodide  solution.  A  dark  brown  liquid 
results.  This  is  known  as  Lugol's  solution,  or  the  "compound  solution 
of  iodine  "  of  the  pharmacopoeia,  when  made  with  certain  definite  quan- 
tities of  iodine,  potassium  iodide  and  water.  Show  that  the  addition  of 
a  great  excess  of  water  produces  a  precipitate  in  this  solution,  and  that 
it  gives  the  blue  color  with  starch. 

Ex.  84.  Dissolve  minute  crystals  of  iodine  in  carbon  disulphide, 
ether  and  chloroform,  and  observe  the  colors  of  the  solutions.  Add 
chlorine  water  to  a  very  dilute  aqueous  solution  of  potassium  iodide  in  a 
test  tube,  until  a  brown  color  is  formed,  and  then  add  several  large 
drops  of  carbon  disulphide  and  shake.  The  iodine  is  taken  from  the 
water  by  the  disulphide,  imparting  to  the  latter  a  characteristic  color. 

Pure  iodine  has  a  specific  gravity  of  4.95.  It  melts 
at  about  115°  and  boils  above  200°. 

Uses.  Iodine  is  employed  in  the  preparation  of  iodo- 
form  and  several  iodides  used  in  medicine.  It  enters  into 
the  composition  of  many  organic  compounds. 


IODINE  AND  HYDROGEN. 

Hydriodic  acid,  HI,  is  a  well-known  and  important 
substance,  best  made  by  the  action  of  phosphorus  and 
water  on  iodine  by  a  process  analogous  to  that  employed 
in  making  hydrobromic  acid.  When  potassium  iodide  is 


GENERAL  CHEMISTRY.  101 

distilled  with  sulphuric  acid,  pure  hydriodic  acid  is  not 
obtained,  as  a  decomposition  of  this  by  the  excess  of  sul- 
phuric acid  follows,  with  liberation  of  iodine. 

In  its  chemical  behavior  this  acid  closely  resembles 
hydrochloric  and  hydrobromic  acxdfy  but/it  is  le^Ss  stable. 
It  decomposes,  liberating  iodjajs.  In  water  it,  is,ex,trernely 
soluble,  yielding  a  heavy  solution?,1,  \  .*->  ;  **<>.'  \\*  \  ,','• 

Two  oxygen  acids  of  iodine  are  known;  one  of  these  is 
called  iodic  acid  and  is  represented  by  the  formula  HIO3. 
It  is  a  white  crystalline  solid,  soluble  in  water,  and  is  best 
made  by  the  oxidation  of  iodine  by  strong  nitric  acid. 
When  strongly  heated  it  decomposes,  yielding  solid  iodine 
pentoxide,  I2O5,  and  this  in  turn  dissolves  in  water  repro- 
ducing the  acid. 

H80  +  I806=2HIOS. 

Several  iodates  are  known;  sodium  iodate  is  found  in 
Chili  saltpeter. 

Finally,  a  more  highly  oxidized  compound  of  iodine  is 
known  and  this  .is  called  periodic  acid,  HIO4.  It  is  a 
colorless  crystalline  solid,  very  soluble  in  water.  Several 
unimportant  compounds  with  chlorine  and  bromine  are 
known,  and  also  a  singular  compound  with  nitrogen  which 
will  be  described  later. 

FLUORINE. 

This  is  a  gaseous  element  of  which  but  little  is  known 
in  the  free  state. 

Occurrence.  It  is  found  in  nature  in  two  important 
mineral  compounds.  One  of  these  is  calcium  fluoride, 
CaF2,  and  is  called  fluorspar.  The  other  is  a  so-called 
double  fluoride,  containing  sodium  and  aluminum,  AlF3-f- 
3NaF,  called  cryolite.  Fluorine  is  found  in  bones  and 
especially  in  the  teeth  in  small  amounts. 

History.  Some  combinations  of  fluorine  have  been 
known  for  many  years,  but  all  attempts  to  isolate  the  ele- 
ment failed  until  quite  recently.  Moissan  succeeded  a  few 


102  GENERAL  CHEMISTRY. 

years  ago  in  liberating  it  by  the  electrolysis  of  hydrofluoric 
acid  in  a  platinum  tube  at  a  very  low  temperature. 

Properties;  Fluorine  is  characterized  by  its  remarka- 
bly strcng'affiniuies.r  Jtc  combines  with  nearly  all  elements 
except  o;&y,gen,,.  It  attacks  »glass  to  combine  with  its  silicon 
and  corrodes  :metais  at-the  ordinary  temperature  quickly. 
Because  of  these  peculiarities  chemists  have  found  difficulty 
in  studying  it  rather  than  in  decomposing  its  compounds. 
At  the  low  temperature  of  Moissan's  experiments  it  may, 
however,  be  set  free  in  platinum.  It  is  a  yellow  gas  which 
decomposes  water  instantly  forming  hydrofluoric  acid  and 
oxygen.  Many  substances  burn  with  the  gas  as  they 
would  with  oxygen.  It  has  recently  been  liquefied. 


FLUORINE  AND  HYDROGEN. 

An  important  combination  of  these  elements  is  known. 
This  compound  is  known  as  hydrofluoric  acid,  HF,  and  is 
very  soluble  in  water.  The  solution  is  now  an  article  of 
commerce  and  is  sold  for  several  purposes.  It  is  usually 
prepared  by  the  action  of  sulphuric  acid  on  calcium 
fluoride,  a  native  mineral  substance  found  in  quantity  in 
several  localities.  This  reaction  is  analogous  to  that  by 
which  hydrochloric  acid  is  made  from  sodium  chloride  by 
means  of  sulphuric  acid,  and  may  be  illustrated  by  this 
equation  : 

2HF 


Calcium  _|_  Sulphuric  —  Calcium    i  Hydrofluoric 
fluoride  '       acid  sulphate    '         acid. 

The  fact  that  hydrofluoric  acid  attacks  glass  may  be 
shown  easily  by  experiment. 

Ex.  85.  In  a  lead  dish,  having  a  diameter  of  5  centimeters  or 
more,  make  a  pasty  mixture  of  strong  sulphuric  acid  and  powdered  cal- 
cium fluoride.  Place  the  dish  on  a  sand-bath  and  by  means  of  splinters 
of  wood  support  over  it  a  square  of  glass,  both  surfaces  of  which  have 
been  covered  with  wax.  In  the  center  of  one  of  the  waxed  surfaces 
scratch  some  letters  or  figures  and  expose  this  surface  to  the  action  of 
the  fumes  which  arise  from  the  dish  when  it  is  heated.  The  temperature 


GENERAL  CHEMISTRY.  103 

must  not  be  allowed  to  get  high  enough  to  melt  the  wax.  After  fifteen 
or  twenty  minutes  remove  the  glass  and  scrape  off  the  layers  of  wax. 
It  will  be  observed  that  at  the  exposed  points  the  figures  or  letters  have 
become  fixed  in  the  glass  by  its  corrosion.  Much  of  our  chemical 
graduated  ware  is  marked  in  this  manner. 

Etching  is  frequently  carried  out  by  immersing  the 
glass  article,  properly  protected  by  wax,  in  an  aqueous 
solution  of  hydrofluoric  acid  or  in  a  mixture  of  dilute  sul- 
phuric acid  and  powdered  fluorspar.  No  heat  is  applied, 
but  a  longer  time  must  be  given  to  complete  the  work. 
This  reaction  is  due  to  the  affinity  of  the  fluorine  for  an 
element  of  the  glass,  the  silicon.  Some  of  this  element  is 
dissolved  out  by  the  hydrofluoric  acid,  forming  silicon 
fluoride,  SiF4. 

The  aqueous  solution  of  the  acid  cannot  be  kept  in  glass 
or  iron  vessels.  It  is  handled  on  the  small  scale  in  bottles 
of  hard  paraffin  or  gutta-percha,  and  in  large  quantities  in 
barrels  coated  with  paraffin. 

Some  of  the  fluorides  are  becoming  important  articles  of 
commerce.  The  native  calcium  fluoride  is  largely  used  as 
a  flux  in  the  smelting  of  iron  ores. 

GENERALITIES. 

The  four  elements,  fluorine,  chlorine,  bromine  and 
iodine,  constitute  a  natural  group,  in  which  a  variation  in 
properties  is  closely  related  to  a  variation  in  atomic 
weight.  Fluorine,  the  lightest  element,  has  the  strongest 
affinity  for  hydrogen  and  all  the  metals,  but  it  forms  no 
combination  with  oxygen.  Iodine,  the  heaviest  of  the 
group,  forms  a  stable  compound  with  oxygen,  while  its 
combinations  with  hydrogen  and  the  metals  are  very  easily 
decomposed.  Fluorine  decomposes  water  immediately, 
iodine  not  at  all.  The  decomposition  by  chlorine  is  much 
slower  than  by  fluorine,  while  by  bromine  it  is  extremely 
slow.  Of  the  three  well  known  elements  in  the  group, 
bromine  stands  between  the  lighter  chlorine  and  heavier 
iodine  in  all  important  properties. 

In  the  following  table  some  of  the  most  important  rela- 
tions of  the  four  elements  just  considered  are  pointed  out 
in  form  suitable  for  easy  comparison  : 


104 


GENERAL  CHEMISTRY, 


F. 

Cl. 

Br. 

I. 

Atomic  weight. 

19.06 

38.12 

35.45 
70.90 
1.33 

Decomposes  it 
in  light. 

HC1 

C12O 
C180S 
C1O2 

HOC1 
HOC1O 
HOClOa 
HOClOs 

79.95 
159.90 
3.18 

Decomposes  it 
very  slowly. 

HBr 

Vone    known. 

HOBr 

HOBrO2 
HOBrOg 

J26.85 
253.70 
4.95 

Does  not 
decompose. 

HI 

IiO, 

HOIO2 
HOIOs 

Molecular  weight 

Liquid  density  

Action  on  water.    .    . 

Decomposes  it 
very   readily. 

HF 

None    known. 

Hydrogen  compounds 

Oxygen  compounds 

Oxygen  acids.  . 

NATURE  OF  ACIDS. 

In  the  foregoing  pages  the  term  acid  has  been  fre- 
quently employed,  and  from  the  experiments  made  or  sug- 
gested the  general  composition  and  properties  of  these 
bodies  have  been  indicated.  Hydrochloric  acid  contains 
chlorine  and  hydrogen,  hydrobromic  acid,  bromine  and 
hydrogen,  hydriodic  acid,  iodine  and  hydrogen,  hydroflu- 
oric acid,  fluorine  and  hydrogen.  That  sulphuric  acid  also 
contains  hydrogen  is  evident  from  many  experiments.  It 
will  be  shown  in  the  next  chapter  that  the  very  common 
and  important  nitric  acid  is  also  a  hydrogen  compound. 
In  general,  it  may  be  said  here,  acids  are  bodies  charac- 
terized by  containing  hydrogen,  which  may  be  readily  re- 
placed by  metals  to  form  a  group  of  compounds  known  as 
salts.  When  zinc  is  dissolved  in  hydrochloric  acid  hydro- 
gen escapes  and  a  salt  called  zinc  chloride  is  produced. 
When  zinc  is  dissolved  in  the  other  acids,  zinc  bromide, 
zinc  iodide,  zinc  sulphate,  etc.,  are  formed,  hydrogen  in 
all  cases  being  set  free. 


GENERAL  CHEMISTRY.  105 

Acids  neutralize  solutions  of  bodies  known  as  alkalies 
and  bases,  forming,  as  before,  salts.  We  had  an  illustra- 
tion of  this  in  the  experiment  in  which  hydrochloric  acid 
was  mixed  with  the  solution  of  caustic  soda.  On  evapora- 
tion, common  salt,  or  sodium  chloride,  was  left.  In  a  simi- 
lar manner  a  mixture  of  hydrobromic  acid  with  caustic 
soda  would  yield  sodium  bromide. 

Acids,  bases  and  salts  comprise  by  far  the  larger  num- 
ber of  substances  considered  in  inorganic  chemistry.  In  a 
following  chapter  the  relations  of  these  bodies  to  each 
other  and  to  certain  allied  substances  will  be  pointed  out. 


CHAPTER  V. 


NITROGEN   AND  THE   ATMOSPHERE. 
GAS  PROBLEMS. 

NITROGEN  is  a  gaseous  element  found  in  the  uncom- 
bined  state  in  the  atmosphere,  and  in  combination 
widely  distributed  through  plant,  animal  and  mineral  sub- 
stances. In  vegetable  tissues  it  is  found  in  all  alkaloids 
and  in  all  the  so-called  proteid  compounds.  It  occurs  in 
the  proteids  of  the  animal  kingdom  also  and  in  many  sub- 
stances produced  by  animals.  In  the  mineral  kingdom  it 
is  found  mainly  in  the  substances  called  nitrates,  of  which 
common  saltpeter  and  Chili  saltpeter  are  the  best  illustra- 
tions. r 

History.  Rutherford,  in  1772,  was  apparently  the  first 
to  suggest  by  actual  experiment  the  presence  in  the  atmos- 
phere of  a  gas  incapable  of  supporting  life  and  combustion. 
Scheele  later,  about  1774,  came  to  the  conclusion  that  the 
atmosphere  must  consist  of  two  distinct  gases,  but  it 
remained  for  Lavoisier,  in  1775,  to  make  a  clear  statement 
of  the  nature  of  the  two  important  gases  in  the  atmosphere. 
The  one  not  supporting  life  or  combustion  he  called  azote, 
while  the  name  nitrogen  was  suggested  later  by  Chaptal. 

Preparation.  We  can  obtain  nitrogen  from  the  air  by 
separating  in  some  manner  the  other  important  element, 
oxygen,  from  it.  As  oxygen  enters  readily  into  combina- 
tion with  many  substances,  while  nitrogen  is  inert,  this  can 
easily  be  done.  In  illustration,  the  following  experiment 
may  be  made  : 


GENERAL  CHEMISTRY.  107 

Ex.  86.  We  may  take  advantage  of  the  reaction  between  phos- 
phorus and  oxygen  to  free  the  nitrogen  from  the  latter  element.  To  this 
end  dry  a  very  small  piece  of  phosphorus,  not  larger  than  a  pea,  and 
enclose  it  in  a  little  cylinder  of  wire  gauze.  Attach  a  piece  of  iron  wire 
to  this  cylinder,  as  a  handle,  and  bend  it  into  a  U  shape  with  one  limb 
longer  than  the  other.  The  gauze  cylinder  is  attached  to  the  shorter 
limb.  Now  ignite  the  phosphorus,  hold  it  over  a  vessel  of  water,  and 
then  depress  a  wide  mouth  bottle  of  about  300  to  400  Cc.  capacity,  over 
the  burning  substance,  so  that  the  mouth  of  the  bottle  dips  beneath  the 
surface  of  the  water.  A  volume  of  air  is  thus  confined  and  exposed  to 
the  action  of  the  burning  phosphorus.  In  a  few  seconds  the  combustion 
is  complete,  when  it  will  be  found  that  the  level  of  the  water  in  the 
bottle  is  above  that  in  the  vessel,  this  being  the  case  because  water  must 
ascend  to  take  the  place  of  the  oxygen,  united  with  the  phosphorus. 
After  the  disappearance  of  the  fumes  of  phosphoric  oxide,  by  solution, 
the  remaining  gas  may  be  tested.  Withdraw  the  wire  gauze,  leaving 
the  mouth  of  the  bottle  still  under  water,  then  close  the  mouth  by  a 
glass  plate  and  bring  it  into  the  upright  position  on  the  table.  Test  the 
gas  in  the  bottle  as  oxygen  was  tested,  using,  however,  a  burning  taper 
or  splinter  in  place  of  one  merely  glowing.  The  flame  will  be  extin- 
guished, showing  the  inert  nature  of  the  gas. 

At  the  beginning  of  the  above  experiment  the  heat  of 
the  combustion  expanded  the  air  in  the  bottle  and  drove 
part  of  it  out  before  it  was  completely  acted  on  by  the 
phosphorus.  It  follows,  therefore,  that  the  gas  volume 
left  does  not  accurately  represent  the  proportion  of  nitro- 
gen in  the  original  air.  It  will  be  shown  later  that  the 
nitrogen  should  amount  to  very  nearly  four-fifths  of  the 
whole. 

The  nitrogen  as  obtained  by  the  above  process  is  never 
quite  pure,  but  by  more  elaborate  methods  it  may  be  secured 
from  the  air  in  practically  pure  condition.  We  may  pro- 
duce it  in  a  pure  state  by  the  decomposition  of  certain 
compcunds  containing  it,  and  one  such  method  will  be 
illustrated  here,  in  which  we  use  ammonium  chloride: 

Ex.  87.  Make  a  mixture  of  powdered  potassium  dichromate  and 
powdered  ammonium  chloride,  using  four  parts  by  weight  of  the  former 
to  one  part  of  the  latter.  With  this  mixture  half  fill  an  iron  gas  pipe  re- 
tort, about  20  Cm.  long  and  1.5  Cm.  in  internal  diameter,  the  arrange- 
ment of  which  is  shown  in  the  next  figure. 

The  pipe  is  closed  with  a  cork  and  delivery  tube,  which  dips  be- 
neath the  surface  of  water  in  a  trough.  On  applying  heat  to  the  retort 
its  contents  decompose  with  liberation  of  nitrogen  gas,  which  passes 
through  the  delivery  tube,  and  may  be  collected  by  displacement  of 
water  in  the  usual  manner.  Fill  several  bottles  with  the  gas  and  test 


108 


GENERAL  CHEMISTRY. 


as  follows:  Remove  the  bottles  with  glass  plates  as  in  other  cases.  Into 
one  bottle  thrust  a  burning  taper  or  piece  of  wood.  Dip  into  a  second 
a  deflagrating  spoon  containing  a  small  piece  of  burning  phosphorus, 
and  into  a  third  dip  a  spoon  with  burning  sulphur.  The  flames  will  be 
extinguished  in  all  cases,  showing  the  very  inert  nature  of  the  gas. 
On  withdrawing  the  spoon  with  the  phosphorus  it  may  reignite  in 
the  air. 

The  reaction  by  which  the  nitrogen  was  secured  in  this 
case  is  somewhat  complex,  but  by  observation  of  certain 
details  it  may  be  understood.  In  the  progress  of  the  de- 
composition it  will  be  noticed  that  vapor  of  water  is  given 
off  along  with  the  nitrogen  gas,  as  it  condenses  in  the  de- 


FIG.    14. 


livery  tube.  At  the  end  of  the  experiment,  after  cooling  the 
tube,  its  contents  may  be  shaken  out  and  examined.  In 
place  of  the  red  substance  taken  we  have  now  a  green  pow- 
der, which  is  found  to  be  partly  insoluble  in  water.  By 
mixing  with  water  in  a  beaker,  stirring  a  few  minutes  and 
filtering,  something  passes  through  the  filter,  leaving  the 
green  substance  undissolved.  By  evaporating  the  filtered 
liquid  we  find  a  residue  of  potassium  chloride,  while  the 
green  substance  the  chemist  recognizes  easily  as  chromium 
oxide.  We  have,  therefore,  produced  in  the  experiment, 
nitrogen  gas,  water,  potassium  chloride  and  chromium 


GENERAL  CHEMISTRY.  109 

oxide.  Careful  investigation  shows  that  these  substances 
are  formed  in  the  proportions  illustrated  by  the  following 
equation  : 

K8Cr2O,  +  2NH4Cl  =    N2    +  2KC1   +4H2O  +Cr2O3 

Potassium      i    Ammonium —  M-.  _J_Potassium_l_    \i7_t_      _l_Chromium 

dichroinate     '       chloride  •  '  •     chloride     '  oxide. 

The  combining  weight  of  the  potassium  dichromate  is 
294.4,  while  that  of  the  double  part  of  ammonium  chlo- 
ride taken  is  107.  We  obtain  28  parts  of  nitrogen,  or  less 
than  one-tenth  the  weight  of  the  dichromate  used.  The 
process  therefore  is  somewhat  expensive. 

Many  compounds  containing  nitrogen  may  be  decom- 
posed, to  liberate  this  substance,  under  proper  conditions. 
In  case  of  the  compound  taken,  the  ammonium  chloride,  it 
is  necessary  to  add  something  to  hold  or  fix  the  elements 
with  which  the  nitrogen  is  here  combined,  and  the  di- 
chromate of  potassium  answers  this  purpose  by  furnishing 
potassium  and  oxygen  to  unite  with  these  elements. 

Nitrogen  has  certain  uses  in  the  arts  at  the  present 
time,  but  large  quantities  of  the  gas  in  the  pure  state  are 
not  required  by  them. 

General  Properties.  Nitrogen  gas  may  be  con- 
densed to  a  liquid  by  application  of  cold  and  pressure. 
The  gas  is  but  slightly  soluble  in  water,  and  is  character- 
ized by  its  extreme  inertness  or  lack  of  positive  properties. 
It  is  neither  combustible  nor  a  supporter  of  combustion, 
but  it  does  unite  directly  with  hot  magnesium.  A  liter  of 
the  pure  gas  weighs  1.257  Gm.  under  standard  conditions. 


THE  ATMOSPHERE. 

The  atmosphere  is  a  mixture  of  oxygen  and  nitrogen, 
essentially,  with  smaller  quantities  of  moisture,  argon  and 
carbon  dioxide,  and  traces  of  other  gases.  The  relation 
between  oxygen  and  nitrogen  has  been  suggested  above, 
and  can  be  shown  by  exact  experiments.  The  amount  of 
oxygen  may  be  readily  determined  by  the  following,  which 


110  GENERAL  CHEMISTRY. 

is  merely  a   modification   of  our  first   experiment  on  the 
separation  of  nitrogen. 

Ex.  88.  Procure  a  glass  tube,  sealed  at  one  end,  having  a  length 
of  about  75  Cm.  and  an  internal  diameter  of  1.5  to  2  Cm.  If  it  is  gradu- 
ated it  will  be  so  much  the  better  for  our  purpose,  but  if  it  is  not  an 
approximate  graduation  may  be  made  as  follows  :  Pour  a  small  meas- 
ured volume  of  water  into  the  tube,  when  held  in  a  vertical  position  and 
mark  its  level  by  means  of  a  close  fitting  rubber  ring  shoved  down  the 
tube.  Then  add  the  same  volume  of  water  and  mark  the  new  level  as 
before,  and  repeat  the  operation  until  the  whole  tube  has  been  divided 
into  equal  small  volumes.  The  rings  should  not  be  displaced  by  ordi- 
nary handling.  Now  pour  about  20  Cc.  of  water  into  the  tube  and 
invert  it  in  a  deep  jar  of  water.  By  means  of  a  clamp  on  a  lamp  stand, 
fasten  the  tube  in  such  a  position  that  the  levels  of  the  water  inside  and 
outside  are  the  same  after  the  tube  has  stood  long  enough  to  have  the 
air  temperature.  The  tube  should  have  a  perfectly  vertical  position, 
Note  the  volume  of  the  air  enclosed  with  reference  to  the  rubber  rings. 
Next  scrape  off  a  piece  of  phosphorus,  weighing  two  or  three  grams,  and 
fasten  it  to  a  thin  iron  wire,  a  meter  and  a  half  in  length,  bent  in  the 
middle  so  that  the  two  halves  are  ctose  together.  Now  depress  the  wire 
with  the  phosphorus  in  the  jar  so  that  the  limb  with  the  phosphorus  is 
brought  beneath  the  opening  of  the  tube.  Then  puil  up  the  free  end  of 
the  wireand  guide  the  phosphorus  with  the  fingers  of  the  other  hand  so 
that  it  enters  the  tube.  By  means  of  the  outside  wire  it  can  now  be 
pulled  up  to  the  top  of  the  tube,  nearly,  and  there  it  should  be  left  about 
24  hours.  In  these  manipulations,  care  must  be  taken  not  to  bring  the 
end  of  the  tube  above  the  surface  of  the  water  in  the  jar,  and  so  change 
the  air  volume  once  read  off.  Great  care  must  be  taken  to  prevent  the 
ignition  of  the  phosphorus  while  handling  it.  This  can  be  avoided  by 
keeping  it  wet.  When  once  up  in  the  tube  there  is  no  further  danger. 
The  phosphorus  undergoes  slow  oxidation  and  gradually  combines  with 
the  oxygen  present.  By  the  end  of  24  hours  the  reaction  will  be  com- 
plete. Then  shove  down  the  wire  and  remove  the  phosphorus  carefully 
and  put  it  back  under  water.  Observe  that  the  water  level  in  the  tube 
is  higher  than  before.  By  means  of  the  clamp,  and  without  touching 
the  tube  with  the  hands,  depress  it  until  the  two  water  levels  are  the 
same  again.  Read  off  the  gas  volume  now  enclosed.  The  decrease  in 
volume  represents  the  oxygen  only. 

Instead  of  inserting  the  phosphorus  by  the  method  described,  the 
student  may  find  it  more  convenient  to  close  the  end  of  the  long  tube, 
after  noting  the  gas  volume,  by  means  of  his  finger,  and  then  carefully 
lift  it  out  of  the  water  in  the  jar  and  dip  it  beneath  the  surface  of  water 
in  a  large  bowl.  The  air  volume  remains  unchanged.  The  tube  can 
now  be  inclined  to  one  side  and  the  phosphorus,  scraped  and  fastened 
to  the  wire  in  the  same  bowl,  can  be  readily  shoved  up  into  the  tube.  If 
the  wire  is  soft  enough  there  will  be  no  difficulty  in  bending  it  around 
the  end  of  the  tube  so  that  the  latter  may  again  be  closed  by  the  finger  and 
brought  back  into  the  jar  and  clamped.  Before  beginning  the  experi- 
ment, a  piece  of  phosphorus  of  the  proper  diameter  should  be  selected. 


GENERAL  CHEMISTRY.  Ill 

To  obtain  an  accurate  result  by  the  above  experiment 
it  is  essential  that  the  temperature  of  the  gas  in  the 
tube  at  the  two  readings  remain  the  same,  and  also  that 
the  pressure  of  the  air  outside,  or  the  barometric  pressure, 
remain  unchanged.  These  conditions  are  practically  never 
attained  and  it  is  therefore  necessary  to  make  certain  cor- 
rections to  compensate  for  these  changes.  By  increase  of 
temperature  gas  volumes  expand,  and  therefore,  if  the 
laboratory  is  warmer  at  the  time  of  the  second  reading 
than  at  the  first,  the  volume  read  off  will  be  high  and  the 
loss  (or  the  amount  of  oxygen)  will  be  made  to  appear  too 
low.  If  the  temperature  at  ,the  second  reading  is  lower 
than  before,  the  residual  volume  will  be  low  and  the 
oxygen  will  thus  be  obtained  too  high.  This  effect  of 
temperature  can  readily  be  observed  by  the  student  by 
grasping  the  tube,  still  clamped  in  position,  in  the  hand. 
The  heat  of  the  body,  thus  communicated  to  it,  is  suffi- 
cient to  make  a  marked  depression*  of  the  water  level  in 
the  inside  of  the  tube.  Before  making  readings  the  tube 
should,  therefore,  be  handled  as  little  as  possible. 

Changes  in  air  pressure  outside  change  the  inner 
volume  also.  An  increase  in  the  air  pressure,  indicated  by 
elevation  of  the  barometer,  is  communicated  through  the 
water  and  decreases  the  gas  volume  in  the  tube.  The 
increased  pressure  forces  the  water  down  in  the  jar  and 
therefore,  because  they  are  in  communication,  up  into  the 
tube.  Following  a  decrease  in  barometric  pressure  the  gas 
volume  in  the  tube  will  expand.  As  preliminary  to  an 
explanation  of  the  calculations  of  these  corrections  let  the 
following  experiment  be  made  : 

Ex.  89.  Suspend  a  thermometer  in  such  a  manner  that  its  bulb 
hangs  wuhin  a  few  centimeters  of  the  middle  part  of  a  graduated  tube 
containing  some  air  as  in  the  last  experiment.  When  the  temperature 
appears  "to  be  constant,  and  the  volume  of  the  gas  therefore  stationary, 
raise  or  depress  the  tube  by  means  of  the  clamp  until  the  water  levels 
inside  and  outside  are  accurately  the  same.  Now,  read  off  the  volume 
of  the  air  as  shown  by  the  graduation,  read  the  thermometer  and  the 
height  of  the  barometer  which  should  hang  in  the  immediate  vicinity  of 
the  other  apparatus.  The  temperature  within  the  tube  is  assumed  to  be 
the  same  as  shown  outside.  With  the  water  levels  the  same  the  pressure 
on  the  gas  within  the  tube  must  be  the  same  as  that  of  the  air,  as 


112  GENERAL  CHEMISTRY. 

measured  by  the  barometer.  With  these  data  at  command  let  the 
student  calculate  the  reduced  volume  at  the  assumed  normal  conditions 
of  0°  C.  and  an  air  pressure  of  760°  Mm.,  by  the  method  explained 
below. 

REDUCTION  OF  GAS  VOLUMES. 

Correction  for  Temperature.  In  comparing  gas 
volumes  it  is  necessary  to  refer  them  to*  some  standard 
temperature,  which  by  common  consent  is  0°  C.  As 
intimated  above,  all  gases  expand  by  increase  in  tem- 
perature, and  practically  at  the  same  rate.  By  rate  of  ex- 
pansion or  coefficient  of  expansion  we  understand  the  fraction 
of  its  volume  at  0°  which  it  increases  for  an  increase  of  1° 
in  temperature.  This  rate  is  the  same  for  an  increase 
from  0°  to  1°  as  it  is  for  an  increase  from  10°  to  11°  or 
from  19°  to  20°  practically.  In  other  words,  it  is  constant, 
or  so  nearly  constant  that  we  assume  it  for  our  purpose. 

Let  us  represent  the  volume  of  a  gas  measured  at  0°  by 
V0  and  the  volume  of  the  same  gas  expanded  to  the  tem- 
perature t°,  by  Vt.  What  is  the  relation  of  Vt  to  V0  ?  It  is 
evident  that  this  equation  must  be  true, 

Vt=V0 -[-increase. 

The  increase  in  volume  is  made  up  of  three  factors,  or 
is  the  product  of  three  factors.  One  of  these  is  the  rate 
of  expansion  defined  above,  which  we  will  call  r.  The 
second  factor  is  the  number  of  degrees  of  temperature 
through  which  the  expansion  takes  place,  and  this  we  call 
/,  while  the  third  factor  is  the  amount  of  V0  itself,  or  in 
other  words  is  the  number  of  units  of  volume  in  V0.  The 
actual  increase  in  volume  for  two  liters  would  evidently  be 
twice  as  great  as  for  one  liter.  Our  equation  therefore 
becomes, 

V(t,  =V(0)+V(0)  XrX/, 
or,  by  a  slight  alteration 

V(t)=V(0)(l+r/). 


GENERAL  CHEMISTRY.  113 

Now,  as  intimated,  the  rate  of  expansion  of  all  gases  is 
nearly  the  same  and  amounts  to  ^r  of  their  volume  at  0°, 
for  each  degree  of  increase.  Therefore,  r=-yfa,  or,  ex- 
pressed decimally  0.00366.  Making  this  substitution  we 
have, 

V(t)=V(0)(l  +  0.00366/). 

This  is  a  fundamental  equation  and  by  transformation 
we  get  the  next  one, 


v 

V  in  t  - 


(0)—  1+0.00366/ 

In  illustration  of  these  equations  assume  that  we  have 
100  Cc.  of  air  at  0°  and  wish  to  know  its  volume  when 
warmed  up  to  25°.  That  is,  we  wish  to  find  V(tJ  or  V20. 


V(25)  =  100  (1  +  0.00366X25) 
=  100  (1+0.0915) 
=  109.  15. 

The  new  volume  is  therefore  109.15  Cc.  Conversely,  if 
we  have  given  this  volume  at  25°  and  wish  to  know  what 
it  becomes  when  cooled  to  0°  we  use  the  formula, 

V(0)__109.15 


1.0915 
=  100  Cc. 

As  gases  contract  below  zero  at  the  same  rate  at  which 
they  expand  above  zero  these  formulas  can  be  used  for 
minus  temperatures  by  change  of  sign. 

Correction  for  Pressure.  We  learn  by  experiment 
that  the  volume  of  a  gas  varies  as  the  pressure  changes, 
but  inversely.  That  is,  if  the  pressure  becomes  doubled 
the  gas  volume  is  contracted  to  one  half.  If  the  pressure 
decreases  to  one  half  the  gas  expands  to  fill  double  the 


114  GENERAL  CHEMISTRY. 

volume.  If  V  represents  the  volume  of  a  certain  gas  at 
the  pressure,  P,  and  V'  the  volume  of  the  same  gas  at  the 
pressure,  P',  it  would  follow  from  the  above  that 

VP  =  V'P'  =  V"P",  etc. 

That  is,  the  product  of  the  volume  and  the  correspond- 
ing pressure  is  a  constant.  We  assume  a  standard  pres- 
sure as  we  do  a  standard  temperature  in  the  measurement 
of  gas  volumes  and  this  is  usually  taken  as  the  average 
pressure  of  the  air  at  the  sea  level,  a  pressure  equivalent 
to  that  of  a  column  of  mercury  760  Mm.  in  height.  If  we 
let  P  =  760  Mm.  and  V  the  normal  volume  then  it  would 
follow  that 

V_V'P' 
760* 

That  is,  the  reduced  or  normal  volume  is  equal  to  the 
product  of  the  observed  volume  and  the  observed  pressure 
divided  by  760  Mm.,  the  standard  pressure.  In  illustra- 
tion, if  we  read  off  a  volume  of  150  Cc.  at  740  Mm.  pres- 
sure, as  shown  by  the  barometer  or  other  pressure  gauge, 
the  reduced  volume  must  be 


=  146.05Cc. 


760 


The  effect  of  changes  of  temperature  and  pressure 
are  independent  of  each  other;  we  can  make,  therefore, 
either  correction  first,  and  on  this  result  calculate  the  other 
correction.  Both  corrections  can  be  introduced  in  one 
formula.  Let  v  represent  the  observed  volume,  h  the 
observed  barometric  height  and  V  the  volume  at  0°  and 
760  Mm.  Then 

y-  v     h 

(1  +  0.00366  /)   760 

This  formula  is  used  where  the  gas  is  measured  under 


GENERAL  CHEMISTRY.  115 

a  pressure  exactly  equivalent  to  that  of  the  barometric 
height,  h.  In  practice  certain  modifications  may  be  neces- 
sary. The  calculation  of  the  results  of  the  last  experiment 
is  a  case  in  point.  The  volume  of  the  gas  is  read  off  over 
water;  it  is  therefore  saturated  with  moisture,  which  also 
exerts  some  pressure,  even  at  low  temperature.  The  air 
pressure,  h>  is  therefore  balanced  by  the  true  gas  pressure, 
which  we  can  call  /,  pJus  the  pressure  of  the  aqueous 
vapor  which  we  can  call  w. 

h—p-{-w,  therefore  p=h — w. 
Our  formula  above  then  becomes 
v  (h — #/) 


V  = 


(1  -f-  0.00366  /)  760 


The  tension  of  water  vapor  is  always  expressed  in  mil- 
limeters of  mercury,  and  can  be  found  from  a  table  in  an 
earlier  chapter.  For  three  common  temperatures  it  may 
be  given  here. 

t  w 

15... 12.7  Mm. 

20 17.4  Mm. 

25 , 23.5  Mm. 

Suppose  now,  that  in  the  last  experiment  we  read  off  a 
volume  of  95.5  Cc.  of  air  saturated  with  moisture,  at  a  tem- 
perature of  20°  C,  while  the  barometer  stood  at  745  Mm. 
The  reduced  dry  volume  would  then  be  found  by  substitu- 
tion as  follows  : 

V-       95.5  (745—17.4)       _69486 
(1+0.00366X20)  760~815.6 

In  practice  it  may  not  be  always  convenient  or  desirable 
to  depress  the  measuring  tube  until  the  levels  inside  and 
outside  are  the  same.  This  is  generally  the  case  when 


116  GENERAL  CHEMISTRY. 

gases  are  measured  and  operated  upon  over  mercury.  A. 
column  of  liquid  stands  up  in  the  tube  and  the  pressure  of 
this,  also,  must  be  brought  into  the  calculation.  The 
pressure  of  this  column,  like  that  of  the  aqueous  vapor, 
must  be  subtracted  from  the  height  of  the  barometer  as 
read  off.  If  this  column  is  mercury,  measure  its  height  in 
millimeters  above  the  level  of  the  mercury  in  the  reservoir 
below  and  call  it  m.  The  general  formula  of  reduction 
then  becomes 

v  (h — m — w). 


V  = 


(1+0.00366  /)  760 


In  measuring  gases  over  mercury  it  is  generally  best  to 
add  a  minute  drop  of  water  to  insure  that  they  are  fully 
saturated  with  moisture. 

Ex.  90.  As  an  exercise  in  these  calculations  let  the  student  meas- 
ure a  gas  volume  over  mercury  in  an  accurately  graduated  tube,  and 
make  all  reductions  necessary  by  the  last  formula.  A  good  barometer 
and  a  thermometer  should  be  mounted  near  the  gas  tube.  The  exercise 
should  be  repeated  until  the  principle  involved  is  perfectly  under- 
stood. 

Other  Air  Tests.  The  amount  of  oxygen  in  the 
atmosphere  can  be  easily  and  quickly  found,  by  absorbing 
it  from  a  measured  air  volume,  by  means  of  an  alkaline 
solution  of  pyrogallol.  When  great  accuracy  is  desired 
rather  elaborate  apparatus  must  be  used,  but  the  method 
may  be  illustrated  by  a  very  simple  experiment. 

Ex.  91.  Select  a  long,  narrow  test-tube,  which  may  be  closed  by 
the  thumb,  and  by  means  of  rubber  rings,  or  otherwise,  divide  it  into 
six  equal  divisions.  Then,  holding  the  tube  vertically,  pour  in  enough 
10  per  cent  solution  of  potassium  hydroxide  to  fill  one  division.  Next 
incline  the  tube,  and  by  means  of  a  knife  blade  introduce  about  half  a 
gram  of  dry  pyrogallol  in  such  a  manner  that  it  will  rest  on  the  side 
of  the  tube,  but  above  the  alkali.  Then  close  the  tube  firmly  by  the 
thumb,  and  shake  it  thoroughly.  After  a  minute,  invert  the  tube,  with 
the  mouth  under  water,  and  remove  the  thumb.  It  will  be  seen  that 
water  enters  to  take  the  place  of  the  oxygen  absorbed  by  the  pyrogallol 
and  alkali.  If  the  tube  is  allowed  to  cool  to  its  original  temperature, 
four  divisions  should  remain  filled  with  gas. 

The    most    accurate    determination  of   the    amount   of 


GENERAL  CHEMISTRY.  117 

oxygen  in  the  atmosphere  is  made  by  aid  of  the  eudiometer 
as  described  in  the  chapter  on  water.  A  volume  of  air  is 
measured  in  the  eudiometer  over  mercury,  the  proper  cor- 
rections being  made.  Hydrogen  is  then  introduced,  more 
than  enough  to  combine  with  the  oxygen  present,  and  the 
new  volume  is  measured.  A  spark  is  then  passed  through 
the  gaseous  mixture  as  described,  and  this  causes  the 
oxygen  to  unite  with  the  requisite  amount  of  hydrogen  to 
form  water.  A  contraction  follows  and  after  a  time  the 
new  volume  is  read  off  and  reduced  to  normal  temperature 
and  pressure.  As  two  volumes  of  hydrogen  unite  with  one 
of  oxygen,  one-third  of  the  loss  noted  is  the  amount  of  the 
latter  originally  present  in  the  reduced  air  volume.  In 
100  volumes  of  dry  air  there  are  21  volumes  of  oxygen. 


OTHER  CONSTITUENTS  OF  THE  ATMOSPHERE. 

The  amount  of  nitrogen  in  dry  air  was  assumed,  until 
quite  recently,  to  amount  to  79  volumes  in  100,  the  traces 
of  other  gases  being  very  minute.  But  in  1895  it  was 
found  by  Rayleigh  and  Ramsay  that  a  gaseous  element 
resembling  nitrogen  is  also  present,  and  this  element  has 
always  been  included  in  what  has  been  measured  as 
nitrogen. 

Argon. 

This  is  the  name  which  was  given  to  the  new  element, 
but  up  to  the  present  time  its  properties  have  not  been 
fully  described.  The  amount  of  argon  in  the  air  appears 
to  be  about  1  per  cent.  It  is  especially  characterized  by 
great  inertness  in  power  of  combination,  and  this  accounts 
for  the  fact  that  it  has  so  long  escaped  recognition.  Argon 
has  been  found  elsewhere  as  well  as  in  the  atmosphere, 
and  more  recently  a  second  gaseo-us  element,  termed 
helium,  has  been  found,  which  frequently  accompanies 
argon.  The  amount  of  this  in  the  atmosphere  is  extremely 
small,  but  it  is  more  abundant  in  the  gases  escaping  from 
certain  springs." 

Argon  is  left  as  a  residue  when  a  large  body  of  air  is 


118  GENERAL  CHEMISTRY. 

passed  over  hot  copper  for  the  removal  of  the  oxygen  and 
repeatedly  over  hot  magnesium  for  the  removal  of  the 
nitrogen.  The  nitrogen  may  also  be  removed  by  combin- 
ing it  with  oxygen  by  aid  of  the  electric  spark  and  over 
alkali  to  absorb  the  acid  products  formed.  The  excess  of 
oxygen  is  afterward  removed  by  the  copper  method  or  by 
pyrogallol,  as  described  above. 

Carbon  Dioxide.  The  air  contains  about  3  volumes 
of  this  gas  in  10,000  under  normal  conditions.  But  in  the 
atmosphere  of  crowded  cities  and  in  buildings  it  may  be 
much  increased.  In  the  streets  it  is  sometimes  present  to 
the  extent  of  6  or  7  volumes  in  10,000,  while  in  crowded 
rooms,  poorly  ventilated,  it  sometimes  reaches  15  volumes 
in  10,000,  or  even  higher.  The  gas  is  produced  by  proc- 
esses of  respiration  and  combustion,  and  is  absorbed  from 
the  air  by  the  growth  of  vegetation.  The  total  amount 
of  this  gas  present  in  the  air  is  enormously  great,  being 
estimated  to  exceed  3,000  billions  of  kilograms. 

Moisture,  The  amount  of  aqueous  vapor  present  in 
the  air  varies  within  wide*  limits,  being  largely  dependent 
on  the  temperature.  A  cubic  meter  of  air,  if  fully  saturated 
at  0°,  contains  4.87  Gm.  of  moisture;  at  10°,  under  the 
same  conditions,  it  contains  9.36  Gm.;  at  15°,  12.75  Gm.; 
at  20°,  17.16  Gm.;  at  25°,  22.84  Gm.,  and  at  30°,  30.1  Gm. 
The  air  is  seldom  fully  saturated,  but  often  holds  75 
per  cent  of  this  amount.  Such  an  atmosphere  is  un- 
pleasantly moist.  If  the  amount  of  moisture  is  below 
50  per  cent  of  that  required  for  saturation  the  air  appears 
dry  to  the  skin.  When  an  atmosphere  nearly  saturated 
with  moisture  is  suddenly  cooled  to  a  temperature  below 
that  for  which  the  water  present  is  sufficient  for  saturation 
a  part  of  this  water  must  precipitate  in  the  form  of  rain. 
On  the  other  hand,  if  the  air  remains  warm  it  may.  hold  a 
very  large  amount  of  moisture  without  precipitating,  and 
through  much  of  our  summer  weather  in  the  United  States 
this  is  often  the  case.  Evaporation  from  the  skin  cannot 
take  place  if  the  air  is  already  saturated  with  moisture,  and 
such  an  atmosphere  we  describe  as  a  "close"  one. 


GENERAL  CHEMISTRY.  119 

The  amounts  of  moisture  and  carbon  dioxide  in  the  air 
are  best  determined  by  aspirating  a  given  volume  of  the 
air  through  a  series  of  weighed  absorbing  tubes.  The 
first  of  these  tubes  are  charged  with  substances  to  take  up 
the  water;  the  following  tubes  contain  something  to  absorb 
the  carbon  dioxide,  usually  caustic  potassa.  The  increase 
of  weight  in  the  tubes,  after  passing  the  measured  air 
volume,  gives  the  amounts  of  the  two  substances  absorbed. 

Ammonia.  Traces  of  ammonia,  NH3,are  found  in  the 
atmosphere  at  all  times,  and  usually  combined  as  am- 
monium carbonate.  The  ammonia  is  mostly  derived  from 
the  decomposition  of  nitrogenous  organic  matter  and 
although  very  small  in  relative  amount  in  the  whole  vol- 
ume of  air  it  is  sufficiently  great  to  be  quite  important.  As 
it  is  carried  down  by  the  rain  it  enters  the  soil  and  there 
serves  as  a  valuable  food  for  growing  plants.  A  large  part 
of  the  nitrogen  taken  up  by  certain  crops  probably  comes 
from  the  ammonia  reaching  the  rootlets  in  this  manner. 

Ozone.  Traces  of  ozone,  peroxide  of  hydrogen,  and 
oxides  of  nitrogen  are  also  often  found  in  the  air.  These 
have  been  referred  to  already. 

Traces  of  sulphurous  oxide  and  other  gases  are  usually 
found  in  the  air  of  cities  and  besides  these,  and  of  great 
importance,  should  be  mentioned  the  small  amount  of 
organic  dust  everywhere  present.  This  dust  consists 
partly  of  living  andpartly  of  dead  matter.  In  the  living  mat- 
ter are  included  numerous  minute  microorganisms  which 
are  active  in  promoting  fermentations  and  putrefaction. 
The  dead  organic  matter  comes  largely  from  the  decay 
and  disintegration  of  animal  and  vegetable  substances. 


CHAPTER  VI. 


COMPOUNDS  OF  NITROGEN. 

/^COMBINATIONS  of  nitrogen  formed  by  direct  union 
^*  with  the  gas  are  rare,  but  a  large  number  of  sub- 
stances containing  nitrogen  can  be  made  indirectly.  Some 
of  these  are  of  great  value  and  importance. 


NITROGEN  AND  OXYGEN. 

Nitrogen  combines  with  oxygen  to  form  five  com- 
pounds, which  are  named  as  follows: 

Nitrogen  monoxide  N2O  gas. 

Nitrogen  dioxide       NO  or  N2O2  gas. 

Nitrogen  trioxide  N2O3  volatile  liquid. 

Nitrogen  tetroxide  NO2  or  N2O4  volatile  liquid. 

Nitrogen  pentoxide  N2O6  solid. 

The  production  of  the  first  and  second  of  these  will  be, 
shown  by  experiment. 

Nitrogen  Monoxide.  This  substance  is  known  as 
laughing  gas,  or  sometimes  as  nitrous  oxide,  and  is  readily 
made  by  the  decomposition  of  a  common  crystalline  sub- 
stance known  as  ammonium  nitrate,  as  shown  below.  It 
is  made  on  the  large  scale  at  the  present  time  and  is  sold 
compressed  in  cylinders. 

Ex.  92.  Dry  a  small  flask  holding  about  250  Cc.,  and  fit  it  with  a 
perforated  cork  and  long  delivery  tube  bent  so  as  to  lead  down  from  the 
flask,  when  mounted  on  a  sand-bath,  to  a  trough  of  water.  Pour  into 
the  flask  10  to  15  Gm.  of  dry  ammonium  nitrate,  insert  the  cork  holding 
the  tube,  and  support  the  flask  on  a  sand-bath.  Now  apply  heat,  very 


GENERAL  CHEMISTRY.  121 

gently  at  first,  which  will  soon  melt  the  solid  ammonium  nitrate.  Later, 
gas  bubbles  will  escape  from  it,  and  it  will  appear  to  boil.  Bring  the 
lower  end  of  the  delivery  tube  beneath  the  surface  of  water  in  the 
pneumatic  trough  or  basin,  and  collect  several  bottles  of  the  escaping 
gas  in  the  usual  manner.  As  the  bottles  fill,  remove  them  by  aid  of 
glass  plates,  and  stand  them  on  the  table  in  the  upright  position.  Test 
the  gas  by  burning  in  it  a  splinter  of  wood,  some  charcoal,  sulphur  and 
phosphorus.  For  the  last  two  use  a  deflagrating  spoon.  These  sub- 
stances will  burn  almost  as  well  as  in  oxygen.  In  performing  this 
experiment  observe  certain  precautions.  The  delivery  tube  should  be 
wide,  the  heat  should  not  be  allowed  to  become  higher  than  necessary 
to  decompose  the  substance,  and  before  removing  the  lamp  the  delivery 
tube  should  be  withdrawn  from  the  water  for  reasons  already  explained. 

The  decomposition  of  the  ammonium  nitrate  takes 
place  according  to  this  equation  : 

NH4NO3=N2O  +  2H2O 

Ammonium     Nitrousi     \yater 

nitrate  ~  oxide 

80         =    44      -f     36. 

100  Gm.  of  ammonium  nitrate  yield,  therefore,  55  Gm. 
of  the  oxide.  One  liter  of  the  gas,  at  standard  tem- 
perature and  pressure,  weighs  1.98  Gm.,  from  which  it 
follows  that  100  Gm.  of  the  nitrate  will  yield  nearly  28 
liters  of  the  gas. 

Properties.  At  a  low  temperature  the  gas  may  be 
readily  compressed  to  a  liquid.  It  is  slightly  soluble  in 
water,  one  volume  of  which  dissolves  about  1.3  volumes 
of  the  gas  at  0°.  It  forms  no  chemical  combination  in 
dissolving.  At  a  moderately  high  temperature  the  gas  is 
decomposed  into  its  constituents,  and  this  accounts  for 
the  fact  that  combustions  follow  so  readily  in  it.  If 
metallic  sodium  is  strongly  heated  with  a  measured  volume 
of  the  gas  over  mercury,  combustion  follows,  and  after  the 
residue  of  nitrogen  cools,  it  will  be  found  to  possess  the 
volume  of  the  original  gas,  from  which  it  follows  that  the 
monoxide  contains  its  own  volume  of  nitrogen  gas. 

Uses.  As  laughing  gas  it  has  been  used  for  many 
years  by  dentists  and  surgeons  for  the  production  of  mild 
anaesthesia.  When  employed  for  this  purpose  it  must  be 


122  GENERAL  CHEMISTRY. 

carefully  washed  by  bubbling  through  water  after  leaving 
the  generator. 

Nitrogen  Dioxide  is  a  substance  formed  in  many 
reactions  in  the  laboratory,  and  usually  where  nitric  acid 
is  decomposed  by  one  of  the  heavy  metals.  Hydrochloric 
and  sulphuric  acids  usually  yield  hydrogen,  in  contact  with 
metals,  but  with  nitric  acid  of  about  1.2  sp.  gr.,  nitrogen 
dioxide  is  liberated.  This  can  take  place  only  through 
complete  decomposition  of  a  part  of  the  acid  as  shown  in 
the  next  experiment. 

Ex.  93.  Arrange  a  bottle  as  for  generating  hydrogen  by  action  of 
dilute  sulphuric  acid  on  zinc,  and  charge  it  with  about  20  Gm.  of  copper 
turnings  and  25  Cc.  of  water.  Pour  in  now  through  the  funnel  tube  25 
to  50  Cc.  of  strong  nitric  acid.  At  first  red  fumes  fill  the  bottle  ;  when 
these  have  been  driven  out  connect  with  a  receiving  bottle  over  water, 
and  collect  as  directed  for  hydrogen.  This  experiment  should  be  tried 
where  there  is  a  good  circulation  of  air,  as  the  gas,  or  rather  the  product 
which  it  forms  with  oxygen,  is  very  irritating.  Collect  several  bottles 
of  the  gas,  then  fill  the  generator  with  water  to  check  the  reaction,  throw 
away  the  acid  thus  diluted,  wash  and  save  any  copper  left.  Test  the 
gas  as  follows.  Plunge  a  burning  splinter  in  one  bottle,  and  some 
burning  sulphur  in  another.  It  will  be  seen  that  it  does  not  support 
combustion.  Remove  the  cover  from  a  third  bottle  of  the  gas  and 
bring  the  mouth  of  a  clean,  dry  bottle  over  it.  Holding  the  bottles  to- 
gether invert  them  so  that  their  contents  will  mix.  Observe  that  red 
fumes  are  formed  as  the  gas  comes  in  contact  with  the  air. 

The  chemical  reaction  in  this  experiment  is  somewhat 
complicated,  but  has  been  determined  by  full  investigations. 
It  appears  from  these  investigations,  which  cannot  be  given 
in  detail  here,  that  the  decomposition  of  the  acid  by  the 
metal  takes  place  in  several  stages.  But  the  final  results 
are  probably  represented  by  this  equation  : 

3Cu+8HNO3  =  3Cu(NO3)2+2NO+4H2O. 

A  few  pages  in  advance,  when  the  properties  of  nitric 
acid  are  consfdered,  some  illustrations  of  the  general  behav- 
ior of  this  acid  with  metals  will  be  given. 

The  gas  is  often  called  nitrogen  dioxide  because  of  the 
fact  that,  for  a  given  weight  of  nitrogen,  it  contains  twice 
as  much  oxygen  as  the  f  rst  oxide  described.  The  older 


GENERAL  CHEMISTRY.  123 

view  of  its  structure  is  represented  by  the  formula,  N2O2. 
It  is  now  well  known  that  this  cannot  be  the  formula. 

Properties.  The  gas  is  very  slightly  soluble  in  water 
and  does  not  combine  with  alkali  solutions.  It  may 
therefore  be  bubbled  through  a  solution  of  sodium 
hydroxide  to  purify  it  in  the  method  of  preparation  given 
above,  which  as  described,  does  not  yield  a  perfectly  pure 
product. 

The  gas  combines  with  oxygen  to  form  the  substance 
NO2  or  N2O4,  which  was  illustrated  in  the  above  experi- 
ment. The  gas  does  not  support  the  combustion  of  wood 
or  sulphur,  but  if  phosphorus,  burning  brightly,  be  plunged 
into  it  active  combustion  follows,  because  in  this  case  the 
initial  temperature  is  high  enough  to  separate  the  oxygen 
from  the  nitrogen. 

Nitrogen  Trioxide  is  not  readily  obtained  in  the  pure 
state,  but  an  illustration  may  be  given  of  its  formation. 

Ex.  94.  Let  the  student  mix  about  10  Cc.  of  strong  nitric  acid 
with  a  gram  of  starch  in  a  small  flask.  This  is  placed  on  a  sand-bath 
in  a  fume  closet  and  slowly  heated.  After  a  time  red  vapors  appear  in 
the  flask  and  the  reaction  soon  becomes  violent.  The  lamp  should  then 
be  removed.  Dense  red  vapors  escape  from  the  flask.  These  consist 
of  nitrogen  trioxide  with  some  tetroxide. 

In  this  experiment  the  starch  decomposes  the  nitric 
acid,  taking  a  part  of  its  oxygen  to  form  several  complex 
bodies.  The  remainder  of  the  nitrogen  appears  as  triox- 
ide mainly.  The  trioxide  dissolves  in  cold  water,  giving 
rise  to  a  body  known  as  nitrous  acid  : 

N203  +  H20  =2HN03 
Nitrogen  _l_    wat«r    —    Nitrous 
trioxide    »  acid. 

Nitrous  Acid  in  the  pure  state  is  not  important  as  it  is 
not  stable.  But  it  combines  with  alkalies,  as  caustic  soda 
or  caustic  potassa,  forming  nitrites,  which  are  very  impor- 
tant substances  for  laboratory  and  manufacturing  uses: 


=  NaN02+H20. 
Nitrites  are  often   made  by   the   reduction  of  nitrates, 


124  GENERAL  CHEMISTRY. 

that  is  by  the  removal  of  oxygen  from  the  latter  salts.  Such 
a  reduction  may  be  effected  by  fusion  with  lead,  as  illus- 
trated by  this  equation  : 


All  nitrites  are  readily  soluble  in  water. 

Nitrogen  Tetroxide.  As  explained  above,  this  body  is 
formed  by  the  union  of  nitrogen  dioxide  with  oxygen.  The 
actual  composition  of  the  substance  varies  with  the  tem- 
perature, consisting  mainly  of  N2O4  at  the  ordinary  work- 
ing temperature  of  25°C.  At  a  few  degrees  below  this 
point  it  may  be  condensed  to  a  liquid.  Nitrogen  tetroxide 
is  decomposed  by  contact  with  water,  forming  nitric  acid 
and  nitrogen  dioxide.  In  presence  of  air,  nitric  acid  is 
finally  the  sole  product. 

In  pure  condition  it  is  best  made  by  the  decomposition 
of  lead  nitrate  by  heat.  The  reaction  is  illustrated  by  this 
equation: 


By  passing  the  gaseous  products  of  the  reaction 
through  a  TJ  tube,  immersed  in  a  cooling  mixture,  the  N2O4 
may  be  condensed  to  a  yellow  liquid  and  thus  separated 
from  the  oxygen.  This  liquid  boils  at  22°,  about. 

Nitrogen  Pentoxide  is  a  laboratory  product  of  no 
importance  in  the  pure  state.  It  is  sometimes  called  nitric 
anhydride,  and  when  dissolved  in  water  forms  nitric  acid, 
as  here  illustrated. 

N205  +  H20=2HN03 


It  is  a  white,  crystalline  product,  and  is  usually  made  by 
extracting  water  from  nitric  acid  by  means  of  phosphoric 
anhydride,  P2C>6-  ^  IS  not  stable  and  breaks  up  into  oxy- 
gen and  nitrogen  tetroxide. 


GENERAL  CHEMISTRY.  125 

NITRIC  ACID. 

Occurrence.  This  acid  is  formed  in  traces  in  nature 
when  electricity  passes  through  moist  air.  In  combined 
condition  it  is  found  as  calcium,  sodium,  or  potassium 
nitrate.  Nitrates  are  produced  in  nature  by  a  number  of 
oxidation  processes,  which  take  place  in  the  soil  and  in 
water  and  which  are  of  the  highest  importance,  as  will  be 
more  fully  explained  some  pages  in  advance. 

History.  The  acid  seems  to  have  been  first  made  by 
distilling  a  mixture  of  saltpeter,  alum  and  blue  vitriol,  and 
this  as  early  as  the  ninth  century.  Later,  in  the  middle  of 
the  seventeenth  century,  it  was  made  by  Glauber  by  a 
process  similar  to  that  still  employed,  that  is,  by  distilling 
saltpeter  with  sulphuric  acid.  The  actual  composition  of 
the  acid  was  a  subject  of  lively  discussion  until  after  the 
days  of  Lavoisier  and  Cavendish.  The  latter  finally  gave 
the  true  explanation  of  its  formation  and  structure. 

Preparation.  As  mentioned,  this  substance  is  pro- 
duced in  small  traces  by  certain  natural  agencies,  but  in 
quantity  it  is  always  made  by  the  decomposition  of  some 
nitrate.  We  have  seen  that  common  salt,  or  sodium 
chloride,  is  decomposed  by  sulphuric  acid,  yielding,  as  one 
product,  hydrochloric  acid.  In  a  similar  manner  saltpeter, 
or  the  nitrate  of  potassium,  is  broken  up  by  distilling  it 
with  sulphuric  acid,  yielding  as  the  important  product 
nitric  acid.  We  can  best  illustrate  this  by  the  following 
experiment : 

Ex.  95.  Arrange  a  glass  retort  with  a  flask  as  a  receiver  as  shown 
by  the  next  figure.  Charge  the  retort  with  25  to  30  Cc.  of  strong  sul- 
phuric acid  and  about  50  Gm.  of  powdered  potassium  nitrate.  Mount 
the  retort  on  a  sand-bath  and  apply  heat,  gently  at  first  and  then  with 
the  full  gas  pressure.  The  contents  of  the  retort  become  thin  and  give 
off  reddish  vapors  which  pass  over  into  the  receiver,  kept  cold  by  water, 
and  condense  to  a  reddish  yellow  liquid;  collect  10  to  15  Cc.  of  the  acid 
liquid  and  apply  tests  to  it  as  follows  :  Transfer  it  to  test-tubes  and  in 
one  dip  a  pine  splinter.  After  a  time  note  the  appearance  of  the  wood 
on  its  withdrawal.  Add  a  little  starch  to  the  same  tube  and  heat 


126 


GENERAL  CHEMISTRY. 


gently  (do  this  in  a  fume  closet).  Observe  the  decomposition  already 
explained.  In  a  second  tube  add  some  copper  turnings  to  the  acid. 
Observe  the  gas  given  off,  and  the  color  of  the  solution  formed.  To  the 
acid  in  a  third  tube  add  some  strong  solution  of  ferrous  sulphate,  green 
vitriol,  and  observe  the  brown  color  formed,  which  disappears  rapidly 
at  first,  but  later  remains  as  a  ring  in  the  middle  of  the  liquid  column. 

The  liquid  residue  in  the  retort  solidifies   on  cooling 
and  then  is  removed  with  some  difficulty.     It  is  therefore 


f^AHKUfi  <9-Qtl. 


FIG.    15. 


better  to  pour  it  through  the  tubulure  of  the  retort,  while 
still  hot,  into  a  porcelain  dish.  The  retort  is  then  allowed 
to  cool  and  what  remains  can  be  readily  washed  out. 

The  substance  in  the  dish  should  be  heated  on  a  sand- 
bath,  in  a  fume  closet,  as  long  as  it  continues  to  give  off 
vapors.  What  is  left  soon  becomes  solid  on  cooling  and 
is  easily  recognized  by  the  chemist  as  potassium  sulphate. 
We  have  therefore  here  a  reaction  in  which  potassium 
nitrate  gives  place  to  potassium  sulphate,  and  free  nitric 


GENERAL  CHEMISTRY.  127 

acid  takes  the  place  of  the  sulphuric.      By  the  use  of  our 
symbols  we  can  illustrate  this  change  as  follows: 

H2SO4-h2KNO3=K2SO4-h2HNO3 

Sulphuric_L_  Potassium  -—Potassium  _j_       Nitric 
acid  nitrate.  sulphate     '         acid. 

Nitric  acid,  it  appears,  is  formed  by  the  decomposition 
of  a  nitrate,  and  conversely,  a  nitrate  can  be  made  by  the 
action  of  nitric  acid  on  certain  other  substances,  in  illus- 
tration of  which  make  the  following  experiment : 

Ex.  96.  Dissolve  about  5  Gm.  of  potassium  carbonate  (pearlash)  in 
a  little  water  and  add  to  it  some  nitric  acid,  a  little  at  a  time.  Effer- 
vescence begins  as  the  liquids  mix.  Add  the  acid  slowly  as  long  as  gas 
escapes  after  thoroughly  stirring  the  solution.  Then  add  a  few  drops 
more  of  the  acid  and  evaporate  the  liquid  resulting  in  a  porcelain  dish 
to  complete  dryness,  stirring  well  at  the  end  of  the  operation.  Compare 
the  solid  substance  left,  in  appearance  and  taste,  with  the  nitrate  used 
in  the  last  experiment. 

From  what  has  been  given  above  it  is  evident  that 
nitric  acid  is  a  very  strong  and  corrosive  substance.  In 
the  pure  form  it  is  not  very  stable,  and  therefore  it  appears 
in  commerce  diluted  with  water.  It  could  not  be  safely 
handled  under  other  conditions. 

On  the  large  scale  the  decomposition  of  the  saltpeter 
is  carried  out  in  iron  retorts  or  boilers  holding  sometimes 
tons  of  the  raw  materials.  Sodium  nitrate  or  Chili  salt- 
peter is  now  commonly  used  instead  of  potassium  nitrate, 
as  it  is  much  cheaper.  The  acid  fumes  which  distill  over 
are  collected  in  a  series  of  well  glazed  earthenware  Woulfe 
bottles,  which  contain  a  little  water.  It  is  not  practicable 
to  make  or  handle  absolute  nitric  acid.  The  commercial 
product  always  contains  some  water.  Several  grades  are 
sold  having  specific  gravities  usually  from  1.38  to  1.45, 
corresponding  to  acids  of  61  to  77  per  cent  strength  ap- 
proximately. A  strong  red  acid  is  found  in  the  market, 
known  as  fuming  nitric  acid.  This  acid  contains  oxides 
of  nitrogen  dissolved. 

Nitric  acid  is  most  characteristically  distinguished  by 
its  power  of  oxidation  or  supporting  a  kind  of  combustion. 
Indeed,  there  are  cases  in  which  it  may  be  made  to  give 


128  GENERAL  CHEMISTRY. 

up  oxygen  directly  to  ordinary  combustible  substances,  as 
illustrated  by  the  following  experiment : 

Ex.  97.  Place  a  small  beaker  on  an  iron  dish  containing  sand. 
Pour  20  Cc.  of  strong  fuming  nitric  acid  into  the  beaker.  Grasp  a  rod 
of  charcoal,  about  5  Cm.  long  and  5  Mm.  in  thickness, with  iron  forceps 
at  one  end,  and  ignite  the  other  end  in  a  gas  flame.  When  it  is  glowing 
brightly  dip  the  burning  end  beneath  the  strong  acid  in  the  beaker  and 
observe  that  the  combustion  continues  with  evolution  of  red  fumes  in 
quantity.  The  oxygen  of  the  acid  is  in  part  taken  by  the  charcoal, 
while  the  rest  is  held  by  the  hydrogen  and  nitrogen.  In  performing 
this  experiment  care  must  be  taken  to  avoid  touching  the  glass  with  the 
hot  charcoal.  As  this  might  break  it  and  spill  the  acid,  the  sand  is 
placed  beneath  the  beaker  to  catch  the  acid  if  this  accident  happens. 
The  experiment  must  be  made  in  a  fume  cjoset,  as  the  gases  given  off 
are  very  offensive. 

In  any  experiment  with  nitric  acid  in  which  red  fumes 
are  abundantly  given  off  we  may  be  certain  that  oxidation 
is  taking  place.  That  is,  a  part  of  the  acid,  at  least,  is 
undergoing  decomposition  by  which  oxygen  is  given  to 
some  other  substance.  In  the  above  case  the  charcoal 
burns  at  the  expense  of  the  oxygen,  so  furnished,  as  well  as 
it  does  by  the  aid  of  atmospheric  oxygen. 

This  property  of  oxidation  is  not  confined  to  the  acid 
alone,  but  is  found  in  marked  degree  in  many  of  the  com- 
binations of  nitric  acid,  or  nitrates,  as  can  be  shown  by  the 
following  experiment  with  potassium  nitrate : 

Ex.  98.  Dry  and  powder  about  10  Gm.  of  potassium  nitrate.  Pour 
it  into  an  ordinary  dry  test-tube,  and  heat  this  carefully  in  the  flame  of 
the  Bunsen  burner.  The  nitrate  soon  begins  to  melt,  and  finally  all  of 
it  becomes  liquid.  Continue  to  heat  it  carefully  in  the  flame,  moving 
the  tube  to  and  fro,  to  evenly  distribute  the  heat.  Then  drop  in  a  small 
piece  of  charcoal  and  observe  that  it  soon  ignites  and  burns  with  a  hiss- 
ing noise.  When  the  first  piece  is  consumed  add  a  second,  which  burns 
in  the  same  manner,  the  tube  being  still  held  in  the  flame.  Then  drop 
in  a  small  fragment  of  sulphur,  which  likewise  burns  at  the  expense  of 
the  oxygen  taken  from  the  nitrate.  In  performing  this  experiment  the 
tube  should  be  heated  over  a  sand-bath  to  catch  the  liquid  saltpeter  in 
case  of  breakage,  which,  however,  need  not  happen  if  care  is  taken. 

The  above  experiment  illustrates  in  a  very  marked 
manner  the  oxygen  furnishing  power  of  a  nitrate.  Niter, 
either  potassium  nitrate  or  sodium  nitrate,  is  used  with 
sulphur  and  charcoal  in  the  production  of  ordinary  gun- 


GENERAL  CHEMISTRY.  129 

powder,  which  contains  the  three  substances  in  certain 
proportions.  Nitric  acid  is  employed  in  the  arts  mainly  on 
account  of  its  oxidizing  properties,  rather  than  because  of 
its  acidity. 

A  few  metals  resist  the  action  of  strong  nitric  acid, 
most  of  them  are  dissolved  forming  nitrates,  while  a  few 
are  converted  into  oxides  without  solution.  This  is  trueof 
tin  and  antimony  as  will  be  shown  later  by  experiment. 

The  behavior  of  nitric  acid  in  dissolving  metals  is  not 
perfectly  clear,  and  several  theories  have  been  advanced  to 
account  for  it.  To  discuss  these  would  be  out  of  place  in 
an  elementary  book.  But  one  view  which  has  long  been 
advocated  may  be  illustrated  by  considering  the  reaction 
between  copper  and  nitric  acid.  It  is  assumed  that  this 
takes  place  in  two  stages.  We  have  first,  apparently, 
what  may  be  called  the  normal  reaction  between  a  metal 
and  an  acid,  that  is,  a  solution  of  the  metal  with  liberation 
of  hydrogen,  as  follows: 

Cu  +  2HNO3  =  Cu(NOs\  +     H2 

Copp«r-(-      NitF*c       =         Copper         J-Hydrogen. 
acid  nitrate 

This  hydrogen,  however,  instead  of  escaping  in  the  free 
state  as  it  does  from  sulphuric  acid  and  zinc,  seems  to 
react  on  the  excess  of  nitric  acid  present  and  decomposes 
it.  Hydrogen  in  this  condition  is  called  nascent  hydrogen, 
and  because  of  its  powerful  attraction  for  oxygen  is  capable 
of  acting  as  a  reducing  agent.  At  any  rate,  it  is  absorbed 
by  the  acid,  and  decomposition  products  of  the  latter 
escape.  The  reaction  here  is  probably  represented  by  this 
equation  : 

3H2   +  2HN03  =  2NO  +4H2O 

Hydrogen+        Nitric        =^ff+  Water. 


By  combining  the  two  equations  we  have  the  one  given 
some  pages  back  : 

3Cu+8HNO3=3Cu(NO3)2+2NO+4H2O. 

According  to  other  views  the  reaction  is  possibly  ex- 
pressed by  these  equations: 


130  GENERAL  CHEMISTRY. 

2HNO3+     3Cu     =  H2O+3CuO-f2NO 
CuO+2HNO3  =  Cu(NO3)2   +  H2O. 

With  certain  metals  and  dilute  nitric  acid  this  assumed 
reduction  by  nascent  hydrogen  seems  to  go  much  further, 
leaving  ammonia  as  the  final  product.  Dilute  nitric  acid 
and  zinc  seem  to  react  on  each  other  in  this  way  : 

4Zn+8HNO3-4Zn(NO3)2+4H2 
4H2  +  HNO3=NH3-f3H2O. 

By  combination  we  have  finally  : 

4Zn+9HNOs=NHs+4Zn(NOs)g+3H8O. 

The  ammonia  remains  combined  with  the  excess  of  acid 
as  ammonium  nitrate. 

Physical  Properties.  Absolute  nitric  acid  has  the 
specific  gravity  of  1.53  at  15°.  It  mixes  with  water  in  all 
proportions  and  decomposes  on  heating.  The  pure  strong 
acid  begins  to  boil  at  about  86°  but  decomposes  into 
water,  oxygen  and  nitrogen  tetroxide.  Under  the  normal 
pressure  an  acid  of  68  per  cent  strength  distills  without 
change.  The  concentrated  acid  does  not  dissolve  iron  and 
may  therefore  be  shipped  in  iron  drums. 

Uses.  Nitric  acid  is  employed  in  many  ways.  A  great 
deal  is  used  in  making  nitro-celluloses  and  nitro-glycerol 
for  explosives.  Much  is  used  also  in  making  nitro-benzene 
and  allied  bodies  employed  in  the  color  industries.  Many 
nitrates  are  made  by  the  solution  of  metals,  carbonates  or 
oxides.  These  nitrates  are,  in  some  cases,  employed  as 
oxidizing  agents.  With  hydrochloric  acid,  nitric  acid  yields 
a  valuable  solution  known  as  aqua  regia,  referred  to  below. 

Nitrates.  The  nitrates  found  in  many  places  in  nature 
are  produced  usually  by  a  series  of  oxidation  processes 
from  animal  or  vegetable  matter  containing  nitrogen.  This 
matter  reaches  the  soil  often  in  the  form  of  animal  waste 


GENERAL  CHEMISTRY.  131 

or  excreta.  The  changes  which  it  there  undergoes  by 
which  its  nitrogen  becomes  combined  with  oxygen  are 
largely  the  result  of  the  action  of  minute  microscopic 
vegetable  cells  known  as  bacteria.  Many  of  these  bacteria 
have  the  power  of  decomposing  nitrogenous  organic  mat- 
ter in  such  a  manner  that  in  presence  of  oxygen  the  nitro- 
gen becomes  united  to  it.  Soils  become  enriched  in  this 
manner,  as  the  nitrates  are  among  the  best  foods  for  grow- 
ing plants.  In  many  cases  it  is  probable  that  the  formation 
of  ammoniacal  compounds  from  more  complex  organic 
substances  precedes  the  oxidation  stage  or  production  of 
nitrites  and  nitrates.  This  is  usually  the  case  in  the  de- 
composition of  sewage  in  streams.  Urea  and  other  organic 
bodies  are  broken  down  through  bacterial  agency  and 
ammonia  results.  In  presence  of  sufficient  air  this  ammo- 
nia later  gives  place  to  nitrites  and  nitrates.  If  this  nitri- 
fication in  the  soil  takes  place  in  the  presence  of  calcium 
bicarbonate,  which  is  a  common  condition,  calcium  nitrate 
results.  This  salt  is  very  soluble  and  may  be  carried 
through  the  soil  to  appear  later  on  the  sides  of  caverns  as 
cave  niter.  The  deposit  in  the  Mammoth  Cave  is  an 
illustration.  Certain  nitrates  often  appear  as  an  efflorescence 
on  the  surface  of  soil  in  hot  countries.  In  India  such 
efflorescence,  mainly  potassium  nitrate,  is  abundant  enough 
to  be  collected  as  an  article  of  commerce.  The  great  beds 
of  sodium  nitrate  in  western  South  America  were  doubtless 
produced  by  the  oxidation  of  decaying  marine  vegetation. 
These  deposits  furnish  a  large  part  of  the  niter  in  com- 
merce to-day. 

Aqua  Regia. 

This  is  a  mixture  of  2  to  3  volumes  of  strong  hydro- 
chloric acid  with  1  volume  of  nitric  acid.  A  decomposi- 
tion takes  place  by  which  two  products,  known  as  nitrosyl 
chloride,  NOC1,  and  nitroxyl  chloride,  NO2C1,  are  formed. 
This  mixture  has  strong  solvent  properties,  as  metals  and 
ores  may  be  dissolved  in  it  which  cannot  be  dissolved  by 
either  nitric  or  hydrochloric  acid  alone.  It  is  therefore  a 
valuable  reagent  in  the  laboratory,  as  will  appear  later. 


132  GENERAL  CHEMISTRY. 


NITROGEN  AND   HYDROGEN.— AMMONIA. 

Although  several  compounds  of  hydrogen  with  nitrogen 
are  known,  only  one  of  them,  ammonia,  is  technically  im- 
portant. This  substance  occurs  in  small  amounts  in  air 
and  water,  but  usually  in  combination  with  something 
else. 

History.  Combinations  of  ammonia  were  known  to 
the  alchemists.  One  of  the  most  important  of  these, 
known  as  sal  ammoniac,  was  obtained  by  many  different 
processes,  and  finally  by  distillation  of  animal  refuse.  The 
carbonate  of  ammonium  resulted  from  this  distillation,  and 
this  on  treatment  with  hydrochloric  acid  furnished  the 
chloride  or  sal  ammoniac.  In  1774  Priestley  found  that 
this  substance  when  distilled  with  lime  yields  a  gas,  ex- 
tremely soluble  in  water  and  which  may  be  easily  decom- 
posed by  the  electric  spark.  The  composition  of  this  gas 
was  determined  by  Davy  and  others.  It  is  represented  by 
the  formula,  NH3. 

Preparation  of  Ammonia.  Ammonia  in  the  free 
gaseous  condition  has  been  made  by  the  direct  combina- 
tion of  its  elements  through  the  aid  of  the  electric  spark. 
But  this  preparation  has  no  practical  importance.  On  the 
large  scale,  and  in  laboratory  experiments,  it  is  best  made 
by  the  decomposition  of  certain  of  its  compounds,  called 
salts  of  ammonium.  This  is  illustrated  by  the  next  experi- 
ment. 

Ex.  99.  Arrange  a  glass  flask  and  Woulfe  bottles  as  in  the  produc- 
tion of  hydrochloric  acid,  some  pages  back.  Each  Woulfe  bottle 
should  contain  about  100  Cc.  of  distilled  water.  In  the  flask  mix 
about  30  Gm.  of  ammonium  chloride  and  the  same  weight  of  slaked  lime. 
Add  enough  water  to  make  a  thick  liquid  mixture  on  shaking.  Close  the 
flask,  make  the  connections  and  apply  a  gentle  heat,  which  may  be 
gradually  increased.  A  gas  is  given  off  from  the  heated  mixture,  which 
passes  over  and  collects  in  the  water  of  the  first  Woulfe  bottle  mainly. 
Some  reaches  the  second  bottle,  and  little  or  none  the  small  flask.  After 
the  application  of  strong  heat  during  half  an  hour,  detach  the  bottles 
and  test  their  contents  as  given  below. 

Ex.  100.  Take  half  a  test-tube  full  of  the  liquid  from  each  of  the 
two  Woulfe  bottles  and  the  small  flask,  and  add  to  each  2  drops  of  an 


GENERAL  CHEMISTRY. 


133 


aqueous  solution  of  methyl  orange  (1:1000).  This  solution  imparts  a 
yellow  color  with  alkalies,  which  is  characteristic.  In  the  liquid  from 
the  first  Woulfe  bottle  the  reaction  should  be  strong,  but  much  weaker 
in  the  others.  Add  dilute  hydrochloric  acid  now  to  each  test-tube,  a 
drop  at  a  time,  and  observe  that  in  the  first  case  many  drops  may  be 
necessary  to  change  the  color  from  yellow  to  pink,  but  that  in  the  sec- 
ond and  third  very  much  less  is  necessary.  Acids  impart  a  pink  color 
to  solutions  of  methyl  orange.  Next  repeat  the  same  experiment,  using 
10  drops  of  a  weak  alcoholic  solution  of  phenol-phthalein(l:1000)  instead 
of  the  methyl  orange.  A  deep  crimson  red  color  is  now  obtained  with 
the  first  test-tube,  and  weaker  shades  with  the  others.  When  hydro- 
chloric acid  is  added  these  colors  disappear.  Alkalies  in  general  give  a 
red  color  with  phenol-phthalein,  but  in  acids  there  is  no  color  reaction. 
These  experiments  prove  the  strong  alkalinity  of  the  ammonia  solution. 


*'1G.   16. 


Ex,  101.  Thoroughly  clean  three  small  porcelain  evaporating 
dishes.  In  one  pour  about  10  Cc.  of  the  ammonia  solution  from  the 
first  Woulfe  bottle.  In  another  take  about  the  same  volume  of  dilute 
hydrochloric  acid,  while  in  the  third  equal  volumes  of  the  ammonia  so- 
lution and  dilute  hydrochloric  acid  are  mixed.  During  the  mixing  a 
great  volume  of  white  fumes  is  produced.  Evaporate  the  three  solutions 
slowly  on  a  sand-bath.  In  the  case  of  the  first  no  residue  will  be  left. 
The  same  is  true  of  the  second,  or  hydrochloric  acid  solution,  showing 
the  complete  volatility  of  both  of  these  products.  But  in  the  third  case, 
with  the  mixture,  we  have  left  a  white  residue,  which  is  identical  with 
the  ammonium  chloride  used  in  the  experiment  on  the  production  of 
ammonia.  It  appears,  therefore,  that  while  ammonia  and  hydrochloric 
acid  are  extremely  volatile,  the  product  of  their  union  is  not,  or  at  any 


134  GENERAL  CHEMISTRY. 

rate,  but  slightly  at  the  temperature  employed.  That  it  is  volatile  at 
a  higher  temperature  is  shown  next,  by  heating  the  dish  containing 
the  residue  more  strongly.  Dense  white  fumes  are  given  off,  leaving 
practically  nothing  in  the  dish.  All  of  these  experiments  must  be  made  in 
the  fume  closet. 

The  above  experiments  illustrate  the  method  by  which 
ammonia  is  obtained  on  the  large  scale.  Crude  ammonium 
chloride,  produced  as  a  by-product  in  the  manufacture  of 
illuminating  gas,  is  distilled  with  slaked  lime,  and  the  gas 
given  off  is  collected  in  water.  An  impure  ammonia  water 
is  thus  made,  which  is  saturated  with  hydrochloric  acid 
yielding  ammonium  chloride,  not  yet  pure,  but  much  better 
than  the  first  crude  substance.  As  this  substance  is  vola- 
tile it  may  be  greatly  purified  by  sublimation.  The  sub- 
limed salt  is  sent  into  commerce  and  used  for  many  pur- 
poses. If  this  sublimed  salt  is  heated  again  with  slaked 
lime  a  very  nearly  pure  ammonia  gas  is  given  off  which  may 
be  absorbed  in  distilled  water,  yielding  a  concentrated  am- 
monia solution.  By  neutralizing  this  with  pure  hydro- 
chloric, nitric  or  sulphuric  acid  we  obtain,  on  evaporation, 
pure  chloride,  nitrate  or  sulphate  of  ammonium. 

In  many  large  chemical  works  at  the  present  time  pure 
ammonia  water  is  obtained  in  one  operation  from  gas 
liquor.  Soft  coal  containsa  little  nitrogen,  and  in  the  man- 
ufacture of  illuminating  gas  by  the  distillation  of  such  coal 
the  nitrogen  becomes  converted  into  ammonia  which  is 
carried  along  with  the  gas  until  a  large  washing  tank, 
called  the  hydraulic  main,  is  reached.  Here  the  very  sol- 
uble ammonia  dissolves  and  combines  with  carbonic 
acid  and  hydrogen  sulphide  from  the  gas,  to  form  carbon- 
ate and  sulphide  of  ammonium.  The  water  in  this  hydraulic 
main  has  to  be  frequently  renewed.  To  recover  the  am- 
monia from  it,  it  is  run  into  large  boilers,  mixed  with 
slaked  lime  and  distilled.  The  lime  decomposes  the  ammo- 
nium salts,  setting  free  ammonia  gas,  which  passes 
through  a  series  of  cooling  pipes  and  small  washing  reser- 
voirs and  is  then  absorbed  in  distilled  water. 


GENERAL  CHEMISTRY.  135 

The  changes  referred  to  above  may  be  represented  by 
equations: 


Ca02H2+2NH4Cl  =  CaCl2+2NH3   +2H2O 

Slaked     _l   Ammonium  —  Calcium_|_A  •    _|_w  , 

lime  chloride     —chloride  ^Ammonia  -(-Water 

NH3+      HC1      =  NH4C1 

Hd™bl 


Ammonia+  ^tric      ^Ammonium 

2NH3+H2S04  =  (NH4)2S04 

Ammonia+Sulphuric    =      Ammonium 

Finally,  we  sometimesi  speak  of  the  solution  of  am- 
monia in  water  as  ammonium  hydroxide,  NH4OH,  which 
is  NH3-{-H2O.  If  we  may  assume  that  this  body  actually 
exists  in  solution,  then  our  reactions  should  be  written 
after  this  manner: 


NH4OH+    HC1      =  NH4C1+H2O 

Ammonium  _|_  Hydrochloric  —  Ammonium_|_  w   . 
hydroxide  ~         acid  "    chloride      rwaier' 


We  have  here  a  behavior  analogous  to  that  by  which 
we  obtained  sodium  chloride  from  hydrochloric  acid  and 
sodium  hydroxide,  or  caustic  soda.  The  solution  of  am- 
monia gas  in  water  possesses  the  properties  of  sodium  and 
potassium  hydroxides  in  a  marked  degree,  as  will  be  shown 
later.  The  term  caustic  ammonia  has  been  sometimes  ap- 
plied to  the  solution. 

Other  important  properties  of  the  gas  remain  to  be 
shown.  For  this  purpose  we  may  decompose  more  am- 
monium chloride  with  slaked  lime,  but  the  gas  may  be 
obtained  more  conveniently  by  heating  the  strong  solution, 
or  commercial  ammonia  water.  The  following  experiment 
will  illustrate  this: 

Ex.  102.     Fit  a  flask  holding  about  300  Cc.  with  a  perforated  stop- 


136  GENERAL  CHEMISTRY. 

per,  through  which  pass  a  straight  glass  tube  about  20  Cm.  in  length. 
Pour  about  50  Cc.  of  the  strongest  ammonia  water  of  the  laboratory 
into  the  flask,  close  it  as  explained,  and  then  support  it  on  a  sand-bath 
in  a  stable  position,  with  a  small  ring  of  a  lamp  stand  around  the 
neck  of  the  flask,  for  instance.  Then  apply  heat.  This  will  cause  the 
ammonia  gas  to  escape  from  the  solution  and  pass  out  through  the  tube. 
When  the  escape  of  gas  becomes  rapid  as  shown  by  the  apparent  boil- 
ing in  the  flask,  hold  a  dry  bottle  over  the  open  end  of  the  tube,  mouth 
downward,  so  that  the  gas  may  enter  and  force  the  air  out.  As  the  gas 
is  but  little  more  than  half  as  heavy  as  air,  it  may  be  collected  readily 
in  this  manner.  The  tube  should  reach  nearly  to  the  bottom  of  the  in- 
verted bottle.  After  a  few  minutes  the  escape  of  the  gas  into  the  air 
will  show  that  the  bottle  is  full.  Lift  it  up  carefully  and  close  the  mouth 
with  a  glass  plate.  It  m-ay  be  then  placed,  mouth  still  downward,  on  a 
table,  while  a  second  bottle  of  the  gas  is  being  collected  in  the  same  man- 
ner. Remove  the  second  when  full,  and  collect  finally  a  third,  giving 
to  this  more  time  than  to  the  others.  Remove  this  bottle,  close  it 
quickly  with  a  glass  plate,  and  bring  the  mouth  of  the  bottle,  still  held 
downward,  under  water  in  a  basin.  Then  remove  the  plate,  when  it 
will  be  noticed  that  the  water  rushes  up  and  nearly  or  quite  fills  the 
bottle.  Had  the  air  been  quite  expelled  in  the  collection  of  the  gas, 
the  water  should  completely  fill  the  bottle,  showing  the  quick  absorp- 
tion of  the  gas  by  the  water.  Apply  a  flame  test  to  one  of  the  other 
bottles.  To  this  end  lift  the  bottle  from  the  table,  mouth  still  down, 
and  insert  a  burning  taper.  It  will  be  extinguished  and  no  flame  will  re- 
main at  the  mouth,  showing  that  the  gas  is  neither  combustible  as  is  hy- 
drogen, nor  a  supporter  of  combustion,  as  oxygen.  With  the  third  bot- 
tle of  the  gas  make  this  experiment:  Lift  it  from  the  plate  and  push 
up  into  it  a  strip  of  perfectly  dry  red  litmus  paper.  The  change  of  the 
color  to  blue  will  not  be  rapid.  Then  moisten  a  piece  of  red  litmus 
paper  in  fresh  water  which  has  not  been  exposed  to  the  ammonia  fumes, 
bring  this  to  the  bottle  and  observe  that  the  change  to  blue  is  immedi- 
ate. The  dry  gas  is  not  an  alkali  but  becomes  so  in  the  presence  of 
moisture. 

Properties.  At  the  temperature  of  0°  C.  and  under 
a  pressure  of  760  Mm.  water  absorbs  about  1,150  times  its 
volume  of  the  gas.  A  cubic  centimeter  of  water  absorbs, 
therefore,  over  a  liter  of  the  gas.  At  30°  nearly  500  volumes 
are  absorbed  by  one  volume  of  water.  It  will  be  seen  in 
what  follows  that  ammonia  solution  is  a  reagent  of  great 
value  in  the  laboratory,  as  it  is  employed  for  many  pur- 
poses. The  gas  is  also  readily  soluble  in  alcohol  and  the 
solution  so  made  has  several  applications. 

At  a  low  temperature  dry  ammonia  gas  may  be  readily 
condensed  to  the  liquid  form  and  is  then  usually  called 
anhydrous  ammonia.  At  a  temperature  of  15°  a  pressure  of 


GENERAL  CHEMISTRY.  137 

about  seven  atmospheres  is  required  for  the  condensation. 
This  liquid  boils  at  a  temperature  of  —38°.  Like  all  con- 
densed gases,  anhydrous  ammonia  absorbs  a  large  amount 
of  heat  on  passing  from  the  liquid  to  the  gaseous  condition, 
and  advantage  is  taken  of  this  fact  in  refrigerating  or  in 
the  production  of  ice.  To  accomplish  this  the  condensed 
ammonia  is  allowed  to  expand  from  strong  storage  tanks 
into  a  system  of  pipes.  These  pipes  may  be  arranged 
around  the  walls  and  ceilings  of  rooms  to  be  cooled,  or  they 
may  be  built  in  more  compact  form  and  immersed  in  a  brine 
reservoir.  In  this  case  the  brine  becomes  cooled  to  a  low 
temperature  and  if  it  is  pumped  into  pipes  it  may  be  made 
to  circulate  through  rooms  or  buildings  where  cooling  is 
desired.  The  cold  brine  reservoir  may  also  be  used  in  the 
production  of  artificial  ice.  It  is  simply  necessary  to 
immerse  tanks  of  distilled  water  in  the  brine  and  allow 
them  to  remain  there  a  day  or  more.  The  water  freezes  to 
a  block  of  clear,  pure  ice,  which  is  easily  removed.  The 
brine  for  this  purpose  must  have  a  temperature  of  — 5°  to 
— 10°  centigrade.  The  expanded  ammonia  is  compressed 
again  by  powerful  pumps  and  so  used  continuously. 

One  liter  of  ammonia  gas  under  standard  conditions 
weighs  0.765  Gm.  The  specific  gravity  referred  to  air  is 
0.589;  it  is  therefore  one  of  the  lightest  gases  known.  It 
does  not  support  combustion  and  can  be  burned  only  with 
difficulty.  The  composition  of  the  gas  may  be  easily  deter- 
mined by  decomposing  it  with  the  electric  spark.  This 
yields  one  volume  of  nitrogen  to  three  volumes  of  hydro- 
gen, which  may  be  shown  in  a  eudiometer  similar  to  the 
one  used  in  the  analysis  of  water. 

Hydroxylamin. 

This  is  a  substance  which  may  be  looked  upon  as  am- 
monia, NH3,  in  which  one  atom  of  hydrogen  is  replaced  by 
the  group,  OH,  called  the  hydroxyl  group.  It  is  an  un- 
stable crystalline  compound  whichdecomposes  very  readily. 
Its  solution  in  water  is  more  stable  and  the  salts,  which 
may  be  compared  with  the  ammonium  salts,  are  easily 
made  and  preserved.  The  hydrochloride,  NH2OH.HC1, 


1&8  GENERAL  CHEMISTRY. 

is  the  most  important.  A  reaction  by  which  it  is  often 
made  depends  on  the  reduction  of  the  gas  NO  by  nascent 
hydrogen. 

2NO+3H2=2NH2OH. 

Hydrazin. 

This  is  a  compound  having  the  formula  N2H4  and  made 
by  decomposition  of  certain  complex  organic  compounds. 
It  is  not  stable  in  the  free  condition,  but  occurs  as  a  sul- 
phate, N2H4H2SO4.  This  sulphate  may  be  decomposed 
by  alkalies  yielding  a  hydrate,  N2H4H2O,  which  is  a  liquid 
resembling  ammonia  in  some  properties.  With  acids  this 
hydrate  yields  well  defined  salts.  Some  organic  derivatives 
of  hydrazin  are  bodies  of  great  practical  importance. 

Hydronitric   Acid. 

This  is  a  peculiar  compound  recently  discovered  and 
studied  having  the  formula  HN3.  It  is  best  made  by  this 
series  of  reactions  :  ammonia  gas  is  led  over  sodium  in  a 
heated  porcelain  tube,  yielding  sodium  amid,  NaNH2: 

Na-fNH3=NaNH2  +  H. 

When  this  reaction  is  complete  a  current  of  dry  nitrous 
oxide  is  passed  through  the  tube,  acting  on  the  amid  in 
this  way  : 


The  sodium  nitride,  NaN3,  distilled  with  dilute  sul- 
phuric acid  yields  free  hydronitric  acid,  or  azoimid.  In 
the  pure  condition  the  acid  is  a  clear  mobile  liquid  which 
explodes  spontaneously  with  great  violence.  In  water  so- 
lution it  is  more  stable.  It  forms  salts  with  most  of  the 
metals,  on  some  of  which  it  has  a  marked  solvent  action, 
even  attacking  gold.  It  is  a  poison  and  destroys  the  skin 
rapidly.  The  odor  of  the  acid  is  extremely  disagreeable. 


GENERAL  CHEMISTRY.  139 

When  inhaled  the  vapor  produces  violent  headache. 
The  salts  are  called  nitrides  and  in  many  respects  they 
resemble  chlorides.  The  structure  of  the  acid  is  probably 

N 
H-N 


NITROGEN  AND  THE  HALOGENS. 

Nitrogen  forms  compounds  with  chlorine,  bromine  and 
iodine,  which  are  all  very  singular  in  this  respect  that  they 
are  explosive  to  a  high  degree.  Of  these  the  iodine  com- 
pound is  most  easily  made  and  with  safety. 

Ex.  103.  In  a  small  porcelain  dish  rub  about  half  a  gram  of 
iodine  to  a  fine  powder  and  cover  it  with  a  few  cubic  centimeters  of 
strong  ammonia  water.  Stir  repeatedly  with  a  glass  rod  and  after  half 
an  hour  wash  the  contents  of  the  dish  into  two  filters.  This  residue 
consists  of  a  dark  powder,  commonly  called  nitrogen  iodide.  Wash  it 
on  the  filters  with  a  little  alcohol  to  remove  any  unchanged  iodine,  and 
then  displace  the  alcohol  by  washing  with  water  several  times.  Remove 
the  filters  from  their  funnels  and  hang  them  up  to  dry.  It  will  be 
found  that  when  the  product  is  perfectly  dry  the  slightest  agitation  is 
sufficient  to  explode  it,  with  a  sharp  report. 

There  has  been  some  uncertainty  regarding  the  com- 
position of  these  explosive  bodies,  and  it  appears  that 
under  different  conditions  different  products  are  obtained. 
What  is  commonly  called  nitrogen  iodide  is  probably  a 
mixture  of  NI3  and  NHI2.  The  chlorine  compound  is  a 
yellowish  liquid  having  the  composition  NC13,  probably. 


FURTHER  THEORETICAL  CONSIDERATIONS. 

In  the  third  chapter  an  outline  of  the  atomic  theory  was 
presented  and  it  will  now  be  in  place  to  introduce  other 
points  of  a  theoretical  nature,  as  the  student  has  become 
familiar  with  the  preparation  and  properties  of  several  new 
and  important  substances. 

The  conception  of  atoms  and  molecules  has  been 
explained,  and  it  was  shown  that  the  combination  of  atoms 


140  GENERAL  CHEMISTRY. 

in  these  molecules  takes  place  in  definite  and  fixed  propor- 
tions. In  all  cases  23  parts  of  sodium  combine  with  35.5 
parts  of  chlorine  ;  39.1  parts  of  potassium  combine  with 
35.5  parts  of  chlorine.  Besides  these  no  other  combina- 
tions of  sodium,  potassium  and  chlorine  are  known.  In 
these  cases  an  atom  of  one  element  combines  with  an  atom 
of  the  other.  An  atom  of  hydrogen  weighing  1  combines 
with  an  atom  of  chlorine  weighing  35.5  and  the  result  is  a 
molecule  of  hydrochloric  acid  with  a  weight  of  36  5  on  our 
arbitrary  scale.  But  when  we  come  to  consider  the  com- 
bination of  oxygen  with  hydrogen  we  find  that  the  amount 
of  the  latter  which  unites  with  the  atomic  weight  or  16 
parts  of  oxygen,  to  form  water,  is  just  twice  as  great  as  the 
amount  which  combines  with  35.5  parts  of  chlorine.  The 
oxygen  atom  appears  to  have,  therefore,  double  the  com- 
bining power  of  the  chlorine  atom.  The  amount  of  hydro- 
gen which  combines  with  one  atom  of  nitrogen  to  form 
ammonia  is  three  times  that  which  combines  with  35.5 
parts  or  one  atom  of  chlorine,  and  the  amount  of  hydrogen 
which  combines  with  one  atom  of  carbon  to  form  a  com- 
pound known  as  marsh  gas,  or  methane,  is  just  four  times 
that  which  will  combine  with  one  atom  of  chlorine.  The 
combining  powers  or  valencies  of  the  ultimate  atoms  are 
therefore  different  and  a  study  of  the  whole  number  of 
atoms  known  shows  that  some  resemble  hydrogen  or  chlo- 
rine, and  these  are  called  univalent;  some  resemble  oxygen 
and  are  called  bivalent;  some  resemble  nitrogen  and  are 
called  trivalent;  some  resemble  carbon  and  are  called  quad- 
rivalent, while  other  atoms  have  still  higher  powers  of 
combination.  Of  the  real  nature  of  this  valence  we  know 
nothing  and  our  methods  of  representing  and  describing  it 
are  at  best  crude.  For  sake  of  simplicity  in  writing  formu- 
las where  the  valencies  of  the  atoms  are  expressed  we  make 
use  of  dashes  as  in  the  following  figures: 

H 

/        !  i 

H— ,      —0—,      — N,      — C— ,    H— 0-H,     H— C— H 

1 


GENERAL  CHEMISTRY.  141 

But  there  are  many  cases  in  which  elements  combine 
with  each  other  in  more  than  one  proportion,  cases  in 
which,  apparently,  there  are  multiple  combining  weights. 
It  will  be  shown  that  two  oxides  of  carbon  may  be  pre- 
pared. One  of  these  is  known  as  carbon  monoxide,  and  the 
other  as  carbon  dioxide.  By  exact  analysis  it  has  been 
shown  that  the  ratio  of  the  carbon  to  the  oxygen  in  the 
first  of  these  is  1  :  1.333,  while  in  the  second  it  is  1  :  2.666. 
That  is,  we  have  for  a  given  weight  of  carbon  twice  as 
much  oxygen  in  one  case  as  in  the  other.  Two  oxides  of 
nitrogen  were  prepared,  and  it  was  explained  that  three 
others  are  known.  An  interesting  fact  was  shown  by  the 
analyses  which  chemists  made  of  these.  The  ratios  of 
nitrogen  to  oxygen  in  the  five  compounds  are  given  in  the 
following  table  : 

N  O 

1st 1 0.5714 

2d 1 1.1428 

3d..  ..!..  ..1.7142 


5th 1 2.8571 

The  proportions  of  oxygen  in  these  compounds  stand  to 
each  other  in  the  relation,  1,  2,  3,  4  and  5.  Two  oxides  of 
sulphur  are  known  also,  in  which  the  amounts  of  sulphur 
present  stand  to  each  other  in  the  exact  relation  of  2  to  3. 

In  investigating  fully  these  three  classes  of  compounds 
it  soon  becomes  apparent  that  the  union  of  oxygen  with 
nitrogen,  carbon  or  sulphur  takes  place  always  as  here  rep- 
resented, and  such  facts  give  the  strongest  support  to  the 
atomic  theory. 

But  analysis  shows  and  the  above  table  illustrates  an- 
other point.  It  appears  that  nitrogen,  under  different  con- 
ditions, has  different  capacities  for  holding  or  combining 
with  oxygen,  and  we  express  this  by  saying  that  the  valency 
of  the  nitrogen  varies. 

Nitrogen  is  not  the  only  element  in  which  the  valency 
or  capacity  for  combination  is  variable.  This  seems  to  be, 
indeed,  the  rule  rather  than  the  exception  and  an  attempt 
is  made  in  the  accompanying  table  to  show  the  variations 
in  valency  in  the  more  important  elements  : 


142 


GENERAL  CHEMISTRY. 


TABLE  OF  VALENCY. 


Name. 

Symbol 

Valence. 

Aluminum  

Al 
Sb 
As 
Ba 
Bi 
B 
Br 
Cd 
Ca 
C 
Cl 
Cr 
Co 
Cu 
F 
Au 
H 
I 
Ir 
Fe 
Pb 
Li 
Mg 
Mn 
Hg 
Mo 
Ni 
N 
O 
P 
Pt 
K 
Se 
Si 
Ag 
Na 
Sr 
S 
Te 
Tl 
Sn 
Ti 
W 
U 
Zn 
Zr 

II 

I 
II 
II 
II 
I 
II 
II 
II 
I 
I 
I 
I 
II 
II 
11 
I 
II 
II 
II 
II 
II 
I 
II 

II 

I 
II 

I 

I 
II 
II 
11 

I 
II 

II 
II 

III 
III 
III 

III 
III 
III 

III 
III 

III 
III 
III 

III 
III 
III 

III 

V 
V 

V 
V 

IV 
V 
IV           VI 
IV 

V 
IV 
IV 
IV 

IV 

IV            VI 
IV 
V 

V 
IV 

IV            VI 
IV 

IV            VI 
IV            VI 

IV 
IV 
IV            VI 
IV            VI 

IV 

VII 
VII 

VII 
(VII) 

Antimony  

Arsenic 

Barium 

Bismuth                        

Boron 

Bromine   

Cadmium  .        

Calcium  

Carbon  

Chlorine 

Chromium  

Cobalt 

CoDDer 

Fluorine  .    . 

Gold  

Hydrogen   

Iodine  

Iridium  

Iron  

Lead  

Lithium 

Magnesium 

Manganese       

Mercury                  

Molybdenum       

Nickel            

Nitrogen  

Oxygen  

Phosphorus                        

Platinum                                . 

Potassium                      .  .  . 

Selenium         .        

Silicon  

Silver  

Sodium  

Strontium 

Sulphur         .                  

Tellurium      ... 

Thallium  

Tin       

Titanium  

Tungsten            .  . 

Zinc                

Zirconium  

GENERAL  CHEMISTRY.  143 

Under  the  head  of  the  compounds  of  chlorine  a  list  of 
acids  formed  by  the  union  of  chlorine  with  hydrogen  and 
oxygen  was  given.  The  names  and  formulas  of  these  acids 
are  here  repeated  and  by  means  of  dashes  the  variation 
in  the  valency  of  the  chlorine  is  indicated. 

H  —  O  —  Cl  hypochlorous  acid. 

H—  O—  Cl=zO  chlorous  acid. 


H—  O—  Cl  chloric  acid. 


H—  O—  C1=O  perchloric  acid. 


In  the  above  formulas  the  hydrogen  has  always  a  val- 
ence of  one  and  the  oxygen  a  valence  of  two.  The  formu- 
las indicate  that  the  oxygen  by  its  double  combining  power 
links  the  hydrogen  to  the  chlorine.  In  the  first  formula 
the  chlorine  has  a  valence  of  one;  in  the  second  of  three; 
in  the  third  of  five;  in  the  fourth  of  seven.  Through  this 
increase  in  valence  the  chlorine  is  able  to  hold  more  and 
more  oxygen,  but,  as  already  suggested,  of  the  nature  of 
this  valence  and  the  reasons  for  its  variations  we  know 
nothing. 

Graphic  Formulas.  Formulas  written  like  the  above 
with  the  symbols  detached  and  separated  by  dashes  are 
called  graphic  or  structural  formulas.  It  is  intended  to 
represent  by  them  the  manner  of  combination  of  the  atoms, 
that  is  to  show  which  are  linked  or  joined  in  the  molecule. 
Formulas  in  which  no  attempt  is  made  to  show  structure 
or  the  mode  of  combination  are  called  empirical  formulas. 
For  economy  of  space  these  are  commonly  employed,  but 
graphic  formulas  are  exceedingly  valuable  and  in  many 
case's  almost  indispensable  in  clearly  representing  reac- 
tions. Their  widest  application  is  found  in  explaining  the 
structure  of  complex  organic  compounds,  but  even  in  rep- 
resenting comparatively  simple  substances  their  use  is 


144  GENERAL  CHEMISTRY. 

apparent.  Sulphuric  acid  is  commonly  represented  by  the 
formula  H2SO4,  but  experiment  shows  that  the  atoms  of 
hydrogen  are  linked  to  the  sulphur  by  aid  of  oxygen.  We 
indicate  this  view  then,  by  the  more  complete  formula: 

H— O—  O=0 

The  nucleus  atom  of  sulphur  is  shown  here  as  having  a 
valence  of  six.  Each  oxygen  atom  has  a  valence  of  two 
and  each  hydrogen  atom  a  valence  of  one.  Such  formulas 
will  be  frequently  used  in  what  follows  but  the  student  must 
remember  that  all  such  attempts  to  indicate  structure  are 
at  best  insufficient. 

Molecules  are  formed  by  aggregations  of  atoms  in  space 
of  three  dimensions,  and  not  in  two  dimensions  as  we  are 
for  convenience  obliged  to  show  them,  and  in  no  case,  as 
yet,  is  our  knowledge  accurate  enough  to  indicate  satisfac- 
torily the  space  relations  of  these  atoms,  beyond  what  is 
suggested  above  for  sulphuric  acid. 

The  acids  referred  to  above  are  qualitatively  alike;  they 
all  contain  hydrogen,  chlorine  and  oxygen.  To  distinguish 
between  them  we  employ  certain  prefixes  and  terminations, 
and  the  student  will  observe  that  the  same  are  used  in 
many  similar  cases.  These  are: 

hypo ous. 

ous. 

ic. 

per ic. 

It  will  be  remembered  that  we  have  nitrous  and  nitric 
acids,  and  later  hyposulphurous,  sulphurous  and  sulphuric 
acids,  hypophosphorous,  phosphorous  and  phosphoric 
acids,  and  others  distinguished  in  the  same  manner  will 
be  mentioned.  The  four  chlorine  acids,  the  sulphur  acids, 
and  the  phosphorus  acids  differ  exactly  in  the  same  man- 
ner, and  that  is  in  the  valencies  of  the  characteristic  ele- 
ments present,  and  consequently  in  the  amount  of  oxygen 
held.  The  acids  having  the  lowest  amount  of  oxygen  are  in 


GENERAL  CHEMISTRY.  145 

all  cases  designated  as  hypo — ous  or — ous  acids,  while  those 
with  more  oxygen  or  with  greater  valency  are  called  — ic 
acids.  The  prefixes  hypo  and  per  (or  hyper)  are  from 
Greek  prepositions  meaning  respectively  less  than,  under, 
below,  and  more  than,  over,  above.  The  terminations  ous  and 
ic  are  arbitrary  indications  of  less  or  greater  valence.  A  hypo 
acid  is  therefore  one  in  which  the  characteristic  element 
(the  chlorine,  sulphur,  phosphorus,  etc.)  has  a  lower  valence 
than  it  has  in  the  ous  acid.  In  the  ous  acid  the  valence  is 
less  than  in  the  ic  acid,  and  in  the  ic  acid  it  is  lower  than 
in  the  per — ic  acid. 

It  has  been  stated  already  that  a  salt  is  formed  by  re- 
placing the  hydrogen  of  an  acid  by  a  metallic  atom.  Salts 
formed  from  acids  containing  but  two  elements,  binary 
acids,  take  the  termination  ide.  Thus,  from  hydrochloric 
acid,  HC1,  we  obtain  chlorides,  NaCl,  KC1,  FeCl2  and 
others.  From  hydrobromic  acid  we  have  bromides,  from 
hydriodic  acid  we  have  iodides.  On  the  other  hand,  if  we 
consider  the  so-called  ternary  acids,  those  with  three  ele- 
ments, as  were  illustrated  above,  we  have,  on  replacement 
of  the  hydrogen,  salts  corresponding  to  the  acids.  The 
nomenclature  of  the  salts  may  be  briefly  indicated  as 
follows  : 

hypo — ous  acids  yield  hypo — ite  salts 

ous  acids  yield  ite  salts 

ic  acids  yield  ate  salts 

per — ic  acids  yield      per — ate  salts 

Thus,  we  have  potassium  chlorate  corresponding  to 
chloric  acid,  and  sodium  hypochlorite  corresponding  to 
hypochlorous  acid.  These  designations  are  arbitrary  but 
the  student  should  make  himself  familiar  with  them  as 
early  as  possible,  as  they  are  of  frequent  occurrence 
throughout  the  book. 


CHAPTER  VII. 


SULPHUR  AND  IT5  COMPOUNDS,  SELENIUM 
AND  TELLURIUM. 

OULPHUR  is  an  element  which  occurs  abundantly  in 
O  nature  in  the  free  state  and  in  many  sulphides  and 
sulphates.  It  is  widely  distributed. 

Preparation.  Crude  native  sulphur  occurs  in  Sicily 
in  large  quantities  and  is  refined  by  very  simple  processes. 
Sulphur  melts  at  a  relatively  low  temperature  and  the  first 
refining  consists  in  melting  it  away  from  the  accompany- 
ing earthy  materials.  The  crude  ore  is  heaped  up  in  a 
large  pit  dug  in  a  hillside  and  ignited.  The  combustion  of 
a  part  of  the  sulphur  furnishes  heat  enough  to  melt  the 
rest  which  collects  at  the  bottom  of  the  pit  and  then 
escapes  through  an  opening  leading  to  a  lower  reservoir. 
The  product  so  obtained  is  not  pure  but  may  be  refined  by 
distillation  from  large  retorts. 

Some  sulphur  is  obtained  by  distillation  of  certain 
sulphide  ores  containing  it,  but  the  amount  so  produced  is 
not  important.  At  a  high  heat  a  sulphide  of  iron  decom- 
poses as  here  represented: 

3FeS2=Fe3S4+S3. 

That  is,  one-third  of  the  total  sulphur  maybe  obtained 
in  the  free  state.  It  will  be  recalled  that  the  dioxide  of 
manganese  may  be  broken  up  in  the  same  manner: 


=  Mn3O4-fO3. 
Vast  quantities  of  sulphur  exist  in  deposits   occurring 


GENERAL  CHEMISTRY.  14? 

in  southern  Louisiana,  at  a  considerable  depth  below  the 
surface  of  the  earth.  After  many  futile  attempts  to  mine 
this,  it  has  recently  been  found  possible  to  bring  it  to  the 
top  of  the  ground  in  this  manner:  Holes  are  bored  down 
to  the  deposit,  and  these  are  doubly  piped.  A  large  pipe 
fills  the  boring,  and  inside  of  this  a  smaller  one  goes  down 
to  a  somewhat  greater  depth.  Hot  water  under  great 
pressure  and  at  a  high  temperature,  about  170°  C,  is 
pumped  down  in  the  space  between  the  pipes.  This  melts 
the  sulphur  in  the  deposit  and  ultimately  forces  it  up 
through  the  inner  pipe.  To  aid  in  maintaining  a  high 
temperature,  a  third  small  pipe  passes  down  through  the 
one  which  conveys  the  sulphur  to  the  surface.  A  current 
of  hot  air  is  forced  down  in  this,  but  the  pressure  on  it  is 
less  than  that  on  the  water.  The  molten  sulphur  on 
reaching  the  surface  is  run  into  shallow  pans  to  cool  and 
solidify.  It  is  nearly  pure,  and  for  most  purposes  needs 
no  refining. 

Properties  of  Sulphur.  Sulphur  appears  in  com- 
merce in  fine  powder  or  "flowers  of  sulphur,"  and  as  roll 
sulphur  or  "brimstone."  Both  varieties  are  insoluble  in 
water,  but  soluble  in  carbon  disulphide,  and  in  several 
other  liquids.  The  solubility  in  carbon  disulphide  may  be 
shown  by  experiment. 

Ex.  104.  Take  about  10  Cc.  of  carbon  disulphide  in  a  test-tube 
and  add  3  or  4  grams  of  the  fine  sulphur.  Shake  the  tube  until  all  dis- 
solves, then  pour  the  solution  into  a  small  beaker,  which  leave  in  a  quiet 
place  for  spontaneous  evaporation.  The  sulphur  separates  in  octahedral 
crystals. 

Sulphur  melts  at  a  temperature  of  about  114°  C.  to  a 
thin  yellow  liquid;  at  a  higher  temperature  it  grows 
darker  and  becomes  viscid,  so  that  it  can  be  poured  only 
with  difficulty.  This  and  other  facts  may  be  readily  shown 
by  trial.  The  boiling  point  is  about  440°.  Sulphur  crys- 
tallizes in  several  forms  which  exhibit  different  physical 
properties.  A  form  which  crystallizes  in  octahedra  is  found 
in  nature.  Similar  rhombic  octahedra  are  obtained  by  crys- 
tallization from  carbon  disulphide  as  explained  above. 
These  crystals  have  a  specific  gravity  of  2.05  at  0°.  When 


148  GENERAL  CHEMISTRY. 

sulphur  is  melted  and  allowed  to  cool  slowly,  a  portion  of 
it  separates  in  long  needles  of  the  monoclinic  system. 
These  have  a  lower  specific  gravity  than  the  octahedral 
variety,  viz. :  1.96.  These  monoclinic  prismatic  needles 
may  be  easily  obtained  by  melting  some  sulphur  in  a  test- 
tube  and  allowing  it  to  cool  a  short  time  until  a  solid  crust 
forms  over  the  top.  When  this  is  broken  through  the  still 
liquid  portion  may  be  poured  out,  leaving  a  crystalline 
mass  attached  to  the  sides  of  the  tube. 

Ex.  105.  Melt  15  to  20  Gm.  of  sulphur  in  a  test-tube  gradually, 
observe  the  changes  in  color  and  degree  of  fluidity.  Above  about  250°  C. 
the  melted  mass  grows  thinner.  After  it  has  become  quite  thin  pour  it 
into  some  water  and  allow  it  to  cool.  A  stringy  plastic  mass  is  obtained 
which  is  elastic  like  crude  rubber.  Set  this  aside  and  allow  it  to 
remain  several  days,  and  then  observe  that  it  has  become  brittle,  or  has 
returned  to  the  common  form. 

At  the  ordinary  temperature  the  affinity  of  sulphur  for 
the  metals  is  slight,  but  at  higher  temperatures  many  com- 
binations may  be  easily  made.  It  has  been  shown  already 
that  sulphur  and  copper  unite  very  readily  when  strongly 
heated.  The  same  reaction  will  now  be  shown  with  iron. 

Ex.  106.  In  a  test-tube  mix  some  flowers  of  sulphur  with  about  an 
equal  weight  of  finely  divided  iron.  Fine  filings  or  powder  are  prefer- 
able, drillings  being  usually  too  coarse.  Heat  the  mixture  slowly  in  the 
lamp  flame.  The  sulphur  melts  and  finally  reaches  a  temperature  at 
which  chemical  union  between  the  two  substances,  accompanied  by 
glowing  of  the  mass,  takes  place.  The  iron  burns  with  the  sulphur  as 
it  does  in  oxygen  and  the  product  is  sulphide  of  iron,  or  ferrous  sul- 
phide. When  the  tube  cools,  break  it,  and  examine  the  contents. 
Powder  some  in  a  mortar  and  observe  the  uniform  dark  color.  Put  some 
small  pieces  in  a  test-tube,  add  some  water  and  then  a  little  hydro- 
chloric acid.  Observe  that  a  gas  with  a  very  disagreeable  odor  is 
given  off. 

Uses  of  Sulphur.  Sulphur  is  employed  in  the  manu- 
facture of  sulphuric  acid,  in  gunpowder  (ordinarily  a  mix- 
ture of  sulphur,  saltpeter  and  charcoal),  in  many  varieties 
of  friction  matches  and  in  the  preparation  of  several  sul- 
phides. It  is  employed  in  considerable  quantity  in  the 
vulcanization  of  rubber  and  in  the  preparation  of  so-called 
hard  rubber.  It  has  also  numerous  minor  applications,  in 
medicine  and  in  the  arts. 


GENERAL  CHEMISTRY. 


149 


SULPHUR    AND    OXYGEN. 

Sulphur  forms  two  important  compounds  with  oxygen, 
one  of  which  has  been  referred  to  before  in  the  experiments 
on  oxygen  gas.  At  a  temperature  sufficiently  elevated  the 
union  takes  place  directly  as  has  been  shown.  The  product 
formed  is  sulphurous  oxide,  SO2. 

Sulphurous  Oxide.  This  oxide  is  found  to  a  small 
extent  in  the  atmosphere  of  cities  where  much  soft  coal  is 
burned.  It  is  also  given  off  in  some  volcanic  gases.  For 


FIG.  17. 

experimental  purposes  it  may  be  obtained  by  a  peculiar 
reaction  in  which  copper  is  made  to  decompose  strong  sul- 
phuric acid,  as  explained  below. 

Ex.  107.  Arrange  apparatus  as  shown  in  the  illustration.  The  flask 
should  hold  400  to  500  Cc.  Put  in  it  15  to  20  Gm.  of  copper  in  small 
pieces  or  turnings  and  add  25  Cc.  of  strong  sulphuric  acid.  The 
Woulfe  bottle  contains  a  small  amount  of  water  to  wash  the  gas  passing 
through  it,  and  the  delivery  tube  from  this  can  be  led  into  another  ves- 
sel of  water,  or  into  a  clean  dry  bottle,  for  collection  of  the  gas.  Apply 


150  GENERAL  CHEMISTRY. 

heat  to  the  flask  and  when  gas  bubbles  begin  to  escape  rapidly  lower 
the  flame  to  avoid  too  violent  a  reaction.  Collect  first  several  bottles  of 
the  gas,  by  displacement  of  air,  which  can  be  easily  done  as  the  gas  is 
over  twice  as  heavy  as  air.  In  collecting  the  gas  cover  the  dry  bottle 
as  well  as  possible  with  a  glass  plate  or  perforated  cardboard.  When 
the  bottle  is  full  a  burning  match  held  at  the  mouth  will  be  extinguished. 
As  the  bottles  fill  cover  them  and  set  aside  for  experiment.  Then  lead 
the  delivery  tube  into  a  bottle  of  water  and  allow  this  to  stand  as  long  as 
gas  is  given  off.  The  water  absorbs  the  gas  and  is  used  below. 

For  technical  purposes  sulphurous  oxide  may  be  made 
by  heating  charcoal  with  strong  sulphuric  acid.  A  reduc- 
tion of  the  acid  takes  place  and  oxide  of  carbon  is  formed 
as  well  as  oxide  of  sulphur: 

C+2H2SO4=CO24-2H2O-{-2SO2. 

For  operations  on  the  large  scale,  such  as  bleaching 
and  fumigation,  the  oxide  is  made  by  burning  sulphur  in 
the  air  or  by  roasting  certain  sulphides  called  pyrites. 

Properties  of  Sulphurous  Oxide.  Of  these,  the  odor 
is  most  characteristic,  while  in  its  chemical  behavior  sev- 
eral marked  peculiarities  may  be  easily  shown.  In  col- 
lecting the  gas  the  fact  that  it  is  not  a  supporter  of  com- 
bustion was  shown.  That  sulphurous  oxide  is  a  good 
bleaching  agent  can  be  shown  by  experiments  on  printed 
cotton  goods. 

At  a  low  temperature  the  gas  may  readily  be  condensed 
to  the  liquid  condition.  The  boiling  point  of  the  liquid  is 
— 8°,  and  in  this  form  it  is  an  article  of  commerce,  being 
employed  in  refrigeration. 

Ex.  108.  Moisten  a  strip  of  calico  and  suspend  it  in  a  bottle  of 
the  gas,  allowing  it  to  remain  half  an  hour.  Many  colors  are  com- 
pletely bleached  in  this  time.  It  must  be  noted,  however,  that  some 
colors  are  not  at  all  acted  upon  by  the  gas. 

Ex.  109.  The  marked  solubility  of  sulphurous  oxide  in  water  may 
be  shown  by  means  of  a  bottle  well  filled  by  the  gas.  Invert  the  bottle, 
closed  by  a  glass  plate,  bring  the  mouth  beneath  the  surface  of  water 
and  then  remove  the  plate.  Water  rushes  up  to  take  the  place  of  the 
dissolved  gas,  as  was  the  case  in  the  experiment  with  ammonia.  Lift 
up  the  bottle  by  means  of  the  plate  and  observe  that  the  water  has  an 
acid  reaction,  as  shown  by  the  litmus  paper  test. 


GENERAL  CHEMISTRY.  151 

At  the  temperature  of  0°  1  volume  of  water  dissolves 
nearly  80  volumes  of  the  gas;  at  20°  1  volume  of  water  dis- 
solves 39.5  volumes  of  the  gas;  one  liter  of  the  gas  weighs 
at  0°  2.89  gm. 

Ex.  HO.  Moisten  a  long  splinter  of  pine  wood  in  strong  nitric  acid 
and  dip  it  into  a  bottle  of  the  gas.  Red  fumes  are  formed  of  oxides  of 
nitrogen,  showing  the  decomposition  of  the  nitric  acid.  In  this  reaction 
the  sulphurous  oxide  is  said  to  act  as  a  reducing  agent.  The  meaning 
of  this  term  will  be  explained  below. 

We  have  remaining  the  solution  of  the  gas  in  water, 
obtained  by  direct  absorption.  Tests  may  be  made  with  it 
as  follows: 

Ex.  III.  Pour  portions  of  about  10  Cc.  each  into  several  test- 
tubes.  To  one  add  some  solution  of  potassium  permanganate,  a  few 
drops  at  a  time.  The  deep  purple  color  of  this  solution  gives  place  to 
a  very  light  pink.  Into  another  portion  pour  a  few  cubic  centimeters 
of  a  dilute  solution  of  potassium  dichromate.  Notice  the  change  of 
color  to  green.  Boil  a  few  cubic  centimeters  of  ferric  chloride  in  a  test- 
tube  and  pour  this,  a  few  drops  at  a  time,  into  the  solution  of  sulphur- 
ous oxide,  heating  the  latter  after  each  addition.  The  yellowish  brown 
color  of  the  iron  solution  changes  to  pale  green  by  this  treatment.  These 
are  all  reduction  actions  again. 

Ex.  112.  Pour  about  10  Cc.  of  the  sulphurous  oxide  solution  into 
a  porcelain  dish,  which  place  on  a  sand-bath  and  heat.  Everything 
evaporates  showing  that  the  product  is  volatile.  The  solution  is  called 
sulphurous  acid,  but,  like  the  ammonium  hydroxide,  is  not  stable.  In 
another  dish  pour  about  25  Cc.  of  the  solution,  add  to  it  some  litmus, 
and  then  dilute  solution  of  caustic  soda  until  the  color  turns  blue.  Then 
restore  the  red  color  by  addition  of  more  sulphurpus  oxide  solution,  put 
the  dish  on  a  sand-bath,  and  evaporate  slowly  to  dryness.  A  white  resi- 
due is  left  which  has  a  sharp  saline  taste,  quite  distinct  from  that  of  the 
caustic  soda.  Now  add  a  little  water  to  the  dish  and  some  dilute  hydro- 
chloric acid.  Effervescence  follows  with  escape  of  gas,  which  the  odor 
shows  is  sulphurous  ©xide.  The  caustic  soda  formed  with  the  solution 
sodium  sulphite,  which  is  decomposed  by  the  hydrochloric  acid,  with 
liberation  of  the  sulphurous  oxide  and  formation  of  sodium  chloride. 

Ex.  113.  Pour  10  Cc.  of  the  sulphurous  oxide  solution  into  a  test- 
tube  and  add  to  it  a  few  drops  of  solution  of  barium  chloride.  A  precipi- 
tate forms  which  consists  of  barium  sulphite,  and  which  dissolves  readily 
by  addition  of  a  little  hydrochloric  acid.  Now  repeat  the  experiment, 
but  add  two  or  three  drops  of  strong  nitric  acid  to  the  sulphurous  oxide 
solution,  boil  a  few  minutes  and  complete  as  before.  A  white  precipi- 
tate forms  here  which  does  not  dissolve  on  addition  of  hydrochloric 
acid.  This  is  barium  sulphate. 


152  GENERAL  CHEMISTRY. 

Leave  the  bottle  containing  the  remainder  of  the  sul- 
phurous oxide  solution  uncorked,  but  away  from  dust,  for 
future  experiments. 

We  must  now  turn  to  a  consideration  of  the  chemical 
changes  involved  in  the  experiments  above.  First,  we  have 
the  reaction  by  which  the  sulphurous  oxide  was  produced. 
When  copper  dissolves  in  hot  sulphuric  acid  copper  sul- 
phate is  formed,  while  water  vapor  and  the  sulphurous 
oxide  are  given  off.  Sulphuric  acid  contains,  as  already 
shown,  hydrogen,  oxygen  and  sulphur  in  the  proportions 
given  by  the  formula  H2SO4.  The  liberation  of  the  sul- 
phurous oxide,  therefore,  involves  the  breaking  up  of  this 
group.  It  is  possible  that  we  have  first  a  reaction  analo- 
gous to  one  which  was  suggested  as  taking  place  between 
copper  and  nitric  acid,  viz.: 

Cu+H2SO4  =  CuSO4  +  H2, 

and  that  the  liberated  hydrogen  acts  here,  as  there,  as  a 
reducing  agent;  that  is,  one  which  decomposes  because 
of  its  power  of  combining  with  oxygen.  The  second  re- 
action, then,  would  be  this  : 

H2     +  H2SO4  =  2H2O+       SO2 

Hydrogen  +    SulaPc^ric    =     Water    -+ 


If  we  consider  these  two  reactions  as  taking  place  to- 
gether we  may  write  : 


It  appears  from  this  last  equation,  which  represents 
the  results  of  experiments  made  quantitatively,  that  one- 
half  of  the  sulphuric  acid  which  takes  part  in  the  reaction 
becomes  combined  to  form  a  sulphate,  while  the  other  half 
is  broken  up.  yielding  water  and  sulphurous  oxide. 

Another  explanation  of  the  reaction  between  copper 
and  sulphuric  acid  has  been  proposed  and  that  is  illus- 
trated by  these  equations  : 

Cu-f-H2SOt=CuO  +  H2SO3, 


GENERAL  CHEMISTRY.  153 

copper  oxide  and  sulphurous  acid  being  formed.     The  first 
dissolves  in  the  excess  of  sulphuric  acid  : 


2SO4=CuSO4 

and  the  second  breaks  up  into  water  and  sulphurous  oxide: 
H2S03=H20+S02. 

The  experiments  made  with  the  sulphurous  oxide  in  the 
condition  of  gas  or  in  solution  are  mainly  illustrative  of  one 
thing,  viz.,  its  reducing  or  oxygen  absorbing  power.  In 
the  experiment  in  which  a  splinter  moistened  with  nitric 
acid  was  dipped  into  a  jar  of  the  gas  it  was  observed  that 
reddish  fumes  were  given  off.  Now,  it  has  been  already 
explained  that  the  appearance  of  these  reddish  fumes  in 
reactions  in  which  nitric  acid  is  concerned  is  indicative  of 
a  decomposition  of  the  acid  by  some  substance  which  can 
take  oxygen.  The  reddish  fumes  come  from  the  residue 
left  after  this  breaking  up  of  the  nitric  acid.  Sulphurous 
oxide,  SO2,  has  a  great  tendency  to  take  up  oxygen  and 
form  a  body  called  sulphuric  oxide,  SO3,  and  especially  in 
presence  of  moisture.  The  nitric  acid  here  furnishes  the 
oxygen  for  the  purpose. 

It  was  shown  that  the  deep  purple  solution  of  potas- 
sium permanganate  is  completely  decolorized  by  the  solu- 
tion of  sulphurous  oxide.  The  reaction  here  is  somewhat 
complicated,  but  is  one  in  which  the  potassium  perman- 
ganate parts  with  oxygen,  converting  the  sulphurous  oxide 
and  water  into  sulphuric  acid.  The  solution  of  the  potas- 
sium dichromate,  used  in  the  same  experiment,  behaves 
quite  in  the  same  manner.  The  potassium  dichromate 
acts  here  as  an  oxidizing  agent.  Because  of  their 
complexity  these  reactions  will  not  be  further  explained 
here,  but  they  are  important  ones  and  will  be  taken  up 
later. 

We  have  next  a  case  which  is  somewhat  simpler.  The 
solution  of  sulphurous  oxide  when  treated  with  nitric  acid 
forms  sulphuric  acid.  We  know  this  because  after  the 
mixing  of  the  liquids  we  are  able  to  prove  the  presence  of 


154  GENERAL  CHEMISTRY. 

sulphuric  acid  by  a  reaction  which  will  later  be  shown  to 
be  a  certain  test  [or  the  acid.  What  the  experiment,  carried 
out  in  detail,  shows  to  take  place  is  illustrated  by  this  equa- 
tion: 

O  +  2HNOs=2H8S04+N803. 


That  is,  an  oxide  of  nitrogen  and  sulphuric  acid  are 
formed. 

In  the  last  experiment  made  it  was  directed  to  allow  a 
solution  of  sulphurous  oxide  in  water  to  stand  exposed  to 
the  air.  This  solution  is  often  considered  as  one  of  sul- 
phurous acid,  formed  by  direct  union  of  the  substances, 
thus  : 

S02-fH20  =  H2S03. 

After  standing  some  days  or  weeks  exposed  to  the  air 
we  observe  that  a  change  has  taken  place  in  the  liquid. 
We  find  that  sulphuric  acid  is  present,  as  we  have  tests  by 
which  we  may  readily  distinguish  sulphuric  acid  from  sul- 
phurous. The  latter  acid  takes  up  oxygen  in  this  manner: 


Sulphuric  acid  is  formed,  therefore,  as  an  oxidation 
product  of  sulphurous  acid.  This  leads  to  the  more  de- 
tailed consideration  of  the  properties  of  sulphuric  acid, 
given  later. 

SULPHUROUS  ACID  AND  SULPHITES. 

As  suggested  above  and  as  indicated  by  experiment, 
the  solution  of  SO2  in  water  may  be  considered  as  sul- 
phurous acid.  As  shown,  the  acid  is  not  stable  but  decom- 
poses by  heat  and  on  exposure  to  the  air.  The  sulphites 
are  more  stable  and  several  are  in  common  use.  The  fol- 
lowing equation  illustrates  their  formation: 

,SO,=Na0SOs-h2H2O. 


When  exposed  to   moist  air   most  of  the  sulphites  be- 
come sulphates  by  absorption  of  oxygen. 


GENERAL  CHEMISTRY,  155 

Uses  of  Sulphurous  Oxide  and  Acid.  It  will  be 
shown  later  that  large  quantities  of  sulphurous  oxide  are 
made  as  a  step  in  the  manufacture  of  sulphuric  acid.  It  is 
also  employed  in  bleaching  silk,  woolen  and  straw  articles 
and  in  the  fumigation  of  buildings.  It  is  used  also  to  pro- 
tect trees  and  vines  from  the  ravages  of  certain  pests. 
Strong  solutions  of  sulphurous  acid  or  of  acid  sulphites 
are  employed  in  washing  barrels  and  tanks  or  vats  used  in 
the  manufacture  and  storage  of  beer.  It  acts  here  to  de- 
stroy ferments,  whose  presence  might  spoil  the  product. 
In  the  laboratory  a  solution  of  sulphurous  acid  is  often 
employed  as  a  reagent. 

Hyposulphurous  Acid.  By  the  action  of  zinc  on  a 
solution  of  sulphurous  oxide  a  peculiar  acid  is  formed 
which  is  the  true  hyposulphurous  acid. 

Zn+H2O+2SO2=:ZnSO3  +  H2SO2 

Zinc  Hyposulphur- 

sulphite  cms  acid. 

The  acid  forms  a  yellow  solution  which  absorbs  oxygen 
readily  and  therefore  acts  as  a  strong  reducing  and  bleach- 
ing agent.  A  corresponding  salt  is  made  by  action  of  zinc 
on  a  solution  of  sodium  acid  sulphite: 

Zn+3NaHS08=NaHS08+Na8SOs+ZnS03+H80. 

The  zinc  and  sodium  sulphites  may  be  crystallized  out 
leaving  a  solution  of  the  acid  hyposulphite  which  is  some- 
times used  in  bleaching.  It  has  been  used  in  the  bleach- 
ing of  syrups  and  other  articles  of  food,  but  for  this  pur- 
pose its  employment  should  be  strongly  condemned. 

Sulphuric  Oxide  or  Sulphur  Trioxide.  This  is  a 
substance  having  the  formula  SO3,  and  is  solid  at  the 
ordinary  temperature.  It  may  be  made  by  several  reac- 
tions, and  readily  when  a  mixture  of  sulphurous  oxide  and 
oxygen  gas  is  passed  over  hot  platinum  sponge.  This 
spongy  material  at  a  moderate  temperature  causes  me  two 
gases  to  combine.  It  may  be  more  easily  made  by  the  dis- 
tillation of  a  liquid  known  as  fuming  sulphuric  acid,  which 


156  GENERAL  CHEMISTRY. 

will  be  described  below.  This  acid  has  the  composition 
H2S2O7,  and  when  distilled  decomposes  forming  H2SO4 
and  SO3,  which  may  be  collected  in  a  cool  and  perfectly 
dry  receiver. 

Properties  of  Sulphur  Trioxide.  The  substance 
made  as  described  above  appears  in  the  form  of  dense, 
white  fumes  which  yield  long,  silky  needles  on  cooling. 
This  crystalline  solid  must  be  preserved  in  sealed  glass 
vessels.  If  brought  in  contact  with  the  air  it  immediately 
attracts  moisture  and  becomes  liquid  sulphuric  acid: 

S03+H2O^H2S04. 

The  trioxide  itself  melts  at  a  temperature  of  16°  to 
form  a  colorless  liquid.  In  dry  form  it  is  without  action  on 
litmus  paper. 

SULPHURIC  ACID. 

History.  This  very  important  acid  has  been  known  for 
many  years,  and  was  first  made  by  the  distillation  of  ef- 
floresced green  vitriol,  hence  the  name,  oil  of  vitriol,  which 
still  clings  to  it.  Reference  to  the  acid  appears  first  in  the 
writings  of  the  Arabian  philosopher,  Geber,  and  later,  but 
very  indefinitely,  in  the  works  of  the  earlier  alchemists.  In 
the  fifteenth  century  Basil  Valentine  described  more  clearly 
the  preparation  of  the  vitriol  and  the  distillation  of  the  same. 
Practically  all  the  acid  used  for  300  years  was  made  by 
that  reaction  which  will  be  referred  to  again,  below. 
About  the  middle  of  last  century  a  process  was  discovered 
by  which  sulphuric  acid  may  be  made  by  the  oxidation  of 
sulphurous  oxide.  This  process  was  developed  in  Eng- 
land, while  the  other  grew  to  importance  in  Germany. 
The  common  sulphuric  acid  used  to  day  is  made  by  a  proc- 
ess which  is  a  development  of  the  crude  attempts  first 
made  in  England  about  150  years  ago. 

Preparation.  The  manufacture  of  sulphuric  acid  on 
the  large  scale  involves  a  number  of  reactions  illustrated 


GENERAL  CHEMISTRY.  157 

by  what  has  been  already  given.  First,  sulphur,  or  a  com- 
pound of  iron  and  sulphur,  known  as  iron  pyrites,  is  burned 
in  furnaces  to  form  SO2  by  the  aid  of  oxygen  from  the  air. 
Then  the  gaseous  sulphurous  oxide,  fumes  of  nitric  acid 
and  steam  are  led  together  into  a  large  lead  lined  chamber 
in  which  they  react  on  each  other  as  illustrated  by  these 
equations: 

SO2  +  HNO3=HO(NO2)SO2. 

Nitroso-sulphuric 
acid. 

This  first  product,  nitroso-sulphuric  acid,  is  decom- 
posed by  steam  with  formation  of  sulphuric  acid  and 
oxides  of  nitrogen: 


A  fresh  quantity  of  sulphurous  oxide  entering  the  lead 
chamber  along  with  air  drawn  in  and  more  steam  combine 
to  produce  a  new  quantity  of  the  nitroso-sulphuric  acid. 


to  be  broken  up  by  steam   as   before,  making  a  continuous 
process. 

The  nitroso  product  may  also  suffer  decomposition  in 
this  manner: 

2HO(NO2)SO2+SO2-f2H2O=:3H2SO4+2NO. 

The  oxide  of  nitrogen,  NO,  is  easily  oxidized  and  on 
this  the  continuity  of  the  process  largely  depends.  It  ap- 
pears, therefore,  theoretically,  that  a  very  small  amount  of 
nitric  acid  is  sufficient  to  convert  an  infinitely  large  amount 
of  sulphurous  acid  into  sulphuric  acid.  In  practice  this  is 
not  quite  true  as  there  is  always  some  loss  of  the  oxidizing 
gases.  Oxygen  is  supplied  by  means  of  a  current  of  air 
drawn  into  the  chambers,  and  it  is  of  course  necessary  to 
remove  the  residual  nitrogen.  In  withdrawing  this,  small 
amounts  of  the  nitrogen  oxides  escape  and  are  lost  to  the 
process.  Several  secondary  reactions  take  place  which 


158 


GENERAL  CHEMISTRY. 


also  occasion  slight  losses.  In  the  annexed  illustration 
the  general  arrangement  of  the  lead  chambers  and  fur- 
naces are  shown. 

Sulphur  or  pyrite  is   burned   in    stoves  at  A,  and  the 


FIG.  18. 

* 

fumes  with  excess  of  air  pass  up  into  the  bottom  of  the 
tower,  E.  In  the  upper  part  of  the  sulphur  burner  a  little 
nitric  acid  is  generated  by  the  action  of  a  small  amount  of 
sulphuric  acid  on  sodium  nitrate  and  this  nitric  acid,  in 
vapor,  passes  along  with  the  air  and  sulphurous  oxide. 


GENERAL  CHEMISTRY.  159 

Steam  is  generated  in  the  boiler,  B,  and  this  is  shown  as 
entering  the  lead  chambers  at  different  points.  As  the 
several  substances  come  together  in  the  first  lead  cham- 
ber the  reactions  given  above  begin  but  they  are  not  com- 
pleted until  three  or  four  are  passed.  In  the  third  or 
largest  chamber  the  principal  part  of  the  combination  is 
completed.  By  time  the  last  chamber  is  reached  the 
sulphurous  oxide  is  all  in  combination,  but  mixed  with  the 
nitrogen  and  oxygen  from  the  air  there  is  always  some  of 
the  oxides  of  nitrogen.  To  save  these  the  whole  gaseous 
residue  is  drawn  by  means  of  a  tall  chimney  or  blower 
down  into  the  bottom  of  the  tower  C,  filled  with  coke  or 
hard  brick;  a  stream  of  strong  sulphuric  acid  trickles  down 
over  this  coke  and  absorbs  the  oxides,  but  allows  the  other 
gases  to  escape.  This  strong  acid  with  its  oxidizing  ab- 
sorbed product  flows  finally  into  a  reservoir,  D,  from  which 
it  is  pumped  to  a  tank  above  the  first  compartment,  E,  into 
which  it  is  discharged  slowly.  Here  it  meets  a  stream  of 
water.  The  dilution  causes  it  to  give  up  the  dissolved 
oxides  of  nitrogen,  because  they  are  not  soluble  in  weak 
sulphuric  acid.  The  liberated  oxides  pass  into  the  first 
lead  chamber  again  with  the  sulphurous  oxide  from  the 
burners  and  thus  remain  in  circulation.  To  hasten  oxida- 
tion a  little  fresh  nitric  acid  is  sometimes  allowed  to  enter 
the  second  chamber  and  flow  over  the  cascade  there  illus- 
trated. The  coke  tower  to  absorb  the  oxides  of  nitrogen 
is  called  a  Gay  Lussac  tower,  and  the  tower  above  the  bur- 
ners where  the  strong  mixture  is  discharged,  is  known  as  a 
Glover  tower. 

The  acid  formed  in  the  chambers  has  a  specific  gravity  of 
about  1.6  and  contains  about  70  per  cent  of  actual  H2SO4. 
This  is  called  chamber  acid  and  is  sold  for  many  purposes. 
A  more  concentrated  acid  is  produced  by  boiling  down  this 
chamber  acid  in  leaden  pans  until  a  product  of  1. 70  to  1.72 
specific  gravity  is  reached;  this  contains  77  to  80  per  cent 
of  acid  and  is  strong  enough  for  most  chemical  decomposi- 
tions. Concentration  cannot  be  carried  beyond  this  in 
lead  because  it  dissolves  in  strong  acid.  The  most  con- 
centrated acid  of  commerce  is  made  by  evaporation  in 
large  glass  globes  or  in  platinum  pans.  The  purest  acid  is 


160  GENERAL  CHEMISTRY. 

distilled  from  platinum  stills,  which  increases  its  cost  con- 
siderably. 

Properties.  Pure  sulphuric  acid  has  a  specific  gravity 
of  1.85  at  0°.  It  is  colorless  and  oily  in  appearance.  When 
heated  it  decomposes  slightly,  and  if  the  temperature  of 
Ebullition,  about  338°,  is  reached  an  acid  of  98.5  per  cent 
strength  is  left.  This  acid  may  be  distilled  without  further 
change.  Commercial  sulphuric  acid  contains  several  im- 
purities and  is  often  brownish  in  color  from  the  presence 
of  traces  of  organic  matter.  It  is  seldom  stronger  than  95 
per  cent,  and  is  often  below  90  per  cent.  The  specific  gravity 
of  this  acid  is  usually  between  1.81  and  1.83  at  15°.  Large 
quantities  of  weaker  acids  are  made  and  sold  for  special 
purposes.  The  so  called  chamber  acid\s  sold  without  con- 
centration, just  as  it  leaves  the  lead  chambers,  and  is  used 
in  several  industries. 

Strong  sulphuric  acid  has  a  remarkable  action  on  water, 
which  is  shown  in  the  next  experiment : 

Ex.  114.  Pour  about  5  Cc.  of  water  into  a  beaker  or  test-tube  and 
add  to  it,  stirring  meanwhile,  about  double  its  volume  of  strong  sul- 
phuric acid.  Pour  the  acid  into  the  ^vater,  not  the  reverse.  It  will  be  no- 
ticed that  the  mixture  becomes  very  hot,  and  that  steam  even  may 
escape.  This  is  a  very  characteristic  behavior  of  the  acid. 

Ex.  115.  Pour  10  Cc.,  about,  of  strong  sulphuric  acid  into  a  small 
beaker,  which  leave  uncovered  but  protected  from  dust  for  several  days. 
Notice  at  the  start  the  depth  to  which  the  acid  fills  the  beaker.  From 
time  to  time  look  at  it  and  observe  that  the  liquid  layer  gradually  grows 
deeper.  If  the  beaker  can  be  allowed  to  stand  some  weeks  a  very 
marked  increase  in  the  volume  of  the  liquid  will  be  seen.  At  the  end  of 
this  long  period  repeat  the  last  experiment  by  pouring  the  acid  into 
water.  A  very  high  temperature  on  mixing  will  not  be  observed  now. 
In  fact,  if  the  acid  can  be  allowed  to  stand  long  enough,  an  increase  in 
temperature  may  be  scarcely  perceptible. 

Another  curious  reaction  depending  on  the  same  affinity 
of  sulphuric  acid  for  water  is  shown  in  the  next  experi- 
ment. In  this  case  a  compound  is  decomposed  and  the 
elements  of  water,  hydrogen  and  oxygen,  abstracted. 

Ex.  116.  Make  a  very  strong  solution  of  cane  sugar  by  dissolving 
about  20  Gm.  in  half  its  weight  of  water,  in  a  beaker  holding  300  Cc.  or 
more.  Pour  into  the  syrup  thus  made  an  equal  volume  of  strong  sul- 


GENERAL  CHEMISTRY.  161 

phuric  acid.  In  a  few  seconds  the  mixture  becomes  hot,  blackens  and 
gives  off  steam.  A  large  volume  of  loose,  finely  divided  carbon  sepa- 
rates and  rises  to  fill  the  beaker,  carried  up  by  the  hot  vapor.  The  sugar 
is  a  combination  of  carbon  with  hydrogen  and  oxygen;  the  last  two  are 
taken  by  the  acid  in  the  form  of  water,  while  the  carbon  is  left  in  the 
free  state.  Ordinary  pine  wood  has  a  composition  very  similar  to  that 
of  sugar.  When  a  splinter  is  dipped  in  strong  sulphuric  acid  it  blackens 
for  the  same  reason. 

The  best  tests  we  have  for  the  recognition  of  sulphuric 
acid  in  the  free  state  depend  on  the  principles  just  illus- 
trated. Sulphuric  acid  in  combination,  that  is,  in  sul- 
phates, may  be  recognized  by  other  tests,  illustrated  below. 

The  behavior  of  strong  sulphuric  acid  when  mixed 
with  salt,  saltpeter  and  many  other  substances  is  also  char- 
acteristic. It  has  been  shown  that  in  the  case  of  salt  a  re- 
action follows  in  which  hydrochloric  acid  and  sodium  sul- 
phate are  formed.  With  saltpeter  we  obtain  nitric  acid 
and  potassium  sulphate,  as  shown.  It  can  readily  be 
proven  that  by  heating  a  substance  known  as  sodium  ace- 
tate with  strong  sulphuric  acid  we  obtain  acetic  acid  and 
sodium  sulphate.  In  general,  it  may  be  said,  the  mem- 
bers of  a  large  and  important  class  of  bodies,  known  as 
salts,  are  decomposed  by  this  acid,  yielding  sulphates  and 
other  acids.  From  chlorides  we  obtain  hydrochloric  acid; 
from  nitrates,  nitric  acid;  from  acetates,  acetic  acid;  from 
phosphates,  phosphoric  acid,  and  so  on.  More  will  be  said 
about  these  salts  later.  We  speak  of  sulphuric  acid  as  a 
strong  acid,  because  it  is  able  to  produce  these  decompo- 
sitions. 

The  behavior  of  sulphuric  acid  with  metals  is  inter- 
esting and  has  been  illustrated  by  several  experiments  al- 
ready. Iron  and-zinc,  and  many  of  the  less  common 
metals,  are  dissolved  in  sulphuric  acid,  especially  when  it 
is  diluted,  with  formation  of  sulphate  and  liberation  of 
hydrogen.  Copper,  mercury  and  some  other  metals  are 
dissolved  by  sulphuric  acid  when  it  is  concentrated  and 
hot  with  formation  of  sulphates  and  liberation  of  sulphur- 
ous oxide  instead  of  hydrogen,  as  shown  in  the  case  of 
copper. 

Uses  of  Sulphuric  Acid.  As  has  been  shown,  the 
manufacture  of  most  of  the  other  common  acids  depends  on 


162  GENERAL  CHEMISTRY. 

the  use  of  sulphuric  acid.  It  is  also  largely  used  in  the 
decomposition  of  common  salt  as  a  step  in  the  manufacture 
of  sodium  carbonate  by  the  Leblanc  process.  It  is  used 
in  the  decomposition  of  phosphate  rocks  in  the  manufac- 
ture of  phosphatic  fertilizers,  which  is  a  very  important 
industry.  In  the  manufacture  of  many  organic  dye-stuffs 
the  use  of  strong  sulphuric  acid  is  necessary,  and,  in  fact, 
it  is  used  in  some  stage  in  hundreds  of  technical  processes. 
The  refining  of  petroleum,  the  production  of  glucose,  the 
recovery  of  ammonia  from  gas  liquor  and  the  manufacture 
of  most  modern  high  explosives,  are  all  industries  in  which 
this  acid  is  practically  necessary.  It  has  become  a  com- 
monplace remark  that  the  industrial  development  of  a 
country  may  well  be  measured  by  the  amount  of  sulphuric 
acid  it  uses. 

Oil  of  Vitriol  or  Pyrosulphuric  Acid.  It  was  said 
at  the  outset  that  sulphuric  acid  was  first  made  by  distilla- 
tion of  green  vitriol.  If  this  substance,  which  has  the 
formula- FeSO4.7H2O,  is  dried  it  loses  most  of  its  water 
and  leaves  a  basic  sulphate  having  the  composition 
Fe2S2O9.  When  this  residue  is  distilled  it  breaks  up  as 
here  represented: 

Fe2S209=Fe203+2S03. 

This  SO3  combines  with  the  small  amount  of  water 
left  in  the  product  and  produces  an  acid  of  the  composition 
H2S2O7.  This  is  the  real  oil  of  vitriol,  pyrosulphuric  acid, 
or  fuming  sulphuric  acid.  It  is  made  practically  by  dis- 
tilling the  dried  vitriol  in  small  well  glazed  earthen 
retorts.  When  exposed  to  the  air  it  gives  off  SO3,  which 
combines  with  moisture  to  produce  white  fumes,  H2SO4 
resulting.  This  very  strong  acid  is  made  in  large  quanti- 
ties for  use  in  the  manufacture  of  organic  products  known 
as  sulphonic  acids.  Some  of  these  acids  are  bodies  of  great 
practical  importance  and  certain  ones  among  them  cannot 
be  made  from  the  weaker  acids.  The  name  Nordhausen 
acid  was  formerly  applied  to  this  strong  acid  from  the 
place  in  Germany  where  most  of  it  was  at  one  time  made. 


GENERAL  CHEMISTRY.  163 

The  Sulphates.  Sulphates  are  salts  formed  from 
sulphuric  acid  by  the  replacement  of  its  hydrogen  by 
metals.  Several  illustrations  of  this  have  been  given. 
They  may  be  made,  also,  by  combining  sulphuric  acid  with 
many  salts,  as  explained  above.  Many  of  these  sulphates 
are  important  bodies  and  most  of  them  are  soluble  in 
water.  A  few  are  not  soluble,  and  on  this  fact  a  method 
for  the  recognition  of  the  whole  group  is  based. 

Ex.  117.  Pour  a  weak  solution  of  sodium  sulphate  into  each  of 
three  test-tubes.  To  one  add  some  solution  of  barium  chloride,  to  the 
second  some  solution  of  strontium  chloride,  and  to  the  third  some  solu- 
tion of  lead  acetate.  In  each  case  a  white  precipitate  forms  which, 
after  a  time,  settles  to  the  bottom  of  the  test-tube;  without  waiting  for 
this,  however,  add  to  each  one  of  three  tubes  some  hydrochloric  acid 
and  shake  the  mixture.  No  change  is  observed,  then  warm  it,  and  still 
no  change  appears  in  the  first  two.  A  partial  solution  may  result  with 
the  third.  We  have  here,  therefore,  precipitates  which  are  insoluble  in 
water  and  hydrochloric  acid  and  which  contain  barium,  lead  and  stron- 
tium as  sulphates.  Of  these  the  barium  sulphate  is  the  most  character- 
istic and  important.  Barium  chloride  used  as  above  is  a  test  for  sul- 
phates, and  in  works  on  analytical  chemistry  it  is  shown  that  when  prop- 
erly employed  it  gives  us  accurate  information  concerning  the  presence 
of  sulphates  in  complex  mixtures  even. 

The  reactions  illustrating  these  precipitations  may  be 
given  here.  First,  with  barium  chloride  and  sodium  sul- 
phate we  have: 

BaCl2-fNa2SO4:=BaSO4-f-2NaCl 

Barium     I       Sodium        _   Barium      I     Sodium 
chloride    '      sulphate     —  sulphate    'chloride. 

Then  with  strontium  chloride  we  have: 

SrCl2+Na2SO4  =  SrSO4  +  2NaCl 

Strontiuml        Sodium     -— Strontium    I      Sodium 
chloride     '       sulphate  sulphate     '     chloride. 

With  lead  acetate  we  have: 
Pb(C2H302)2+Na2S04  =  PbS04+2NaC2H302 

Lead  _!_     Sodium      _      Lead       _j_          Sodium 

acetate  sulphate     "       sulphate      '  acetate. 

Barium '  chloride  solution  with  sulphuric  acid  itself 
gives  the  same  white  precipitate,  but  in  this  case  hydro- 
chloric acid  is  formed: 

BaCl8+H8S04=BaS04+2HCl. 


164  GENERAL  CHEMISTRY. 

*| 

It  is  customary  to  speak  oifree  and  combined  sulphuric 
acid.  The  second  expression,  referring  to  the  sulphates, 
is  not  strictly  accurate  as  the  whole  of  the  acid  is  not  in 
combination.  But  the  characteristic  part  is  and  hence  the 
use  of  the  term.  The  same  applies  to  other  salts,  as  the 
nitrates,  the  chlorides,  and  others.  All  acids  have  hydro- 
gen in  common,  but  with  the  hydrogen  we  have  something 
characteristic  for  each  acid.  This  characteristic  element 
or  group  enters  as  combined  sulphuric,  nitric,  hydrochloric 
or  other  acid  into  the  corresponding  salts. 

Other  Sulphuric  Acids. 

Besides  the  above  several  other  acids  containing  hydro- 
gen, oxygen  and  sulphur  are  known,  free  or  in  combination. 
The  most  important  one  of  these  is  called  thiosulphuric 
acid.  This  exists  in  the  well-known  salt,  sodium  thiosul- 
phate,  Na2S2O3-j-5H2O,  formerly  called  hyposulphite  of 
soda.  This  salt  is  largely  used  by  photographers.  The 
free  acid  is  not  stable.  Solutions  of  thiosulphates  de- 
compose with  liberation  of  sulphurous  oxide  and  pre- 
cipitation of  sulphur  when  mixed  with  dilute  acids.  The 
names  and  formulas  of  other  acids  are  these  : 

Dithionic  acid H0S2O6. 

Trithionic   acid H2S3O6. 

Tetrathionie  acid H2S4O6. 

Pentathionic  acid H2S5O6. 

These  acids  are  rare  and  have  no  practical  importance 
in  the  arts.  A  detailed  description  of  them  is  not  called 
for  in  an  elementary  book. 

Of  somewhat  greater  importance  are  the  persulphates 
and  persulphuric  acid,  recently  discovered.  The  acid  has 
the  formula  HSO4,  or  H2S2O8,  and  is  made  by  the  elec- 
trolysis of  strong  H2SO4  at  a  low  temperature.  This  per- 
sulphuric acid  is  not  stable,  but  decomposes  with  water 
yielding  hydrogen  peroxide  : 


GENERAL  CHEMISTRY.  165 

It  is  therefore  a  strong  oxidizing  agent.  Some  of  the 
persulphates  are  of  technical  importance.  The  potassium 
salt,  KSO4  or  K2S2O8,  is  made  by  electrolysis  of  potas- 
sium acid  sulphate,  KHSO4  at  a  low  temperature.  When 
warmed  with  water  it  decomposes  in  this  way: 


It  is  employed  as  an  oxidizer.  An  oxide  correspond- 
ing has  the  formula  S2O7,  but  is  not  technically  useful  in 
the  pure  state.  It  is  not  stable. 

SULPHUR  AND  HYDROGEN. 

Two  combinations  of  these  elements  are  known,  but 
only  one  of  them  is  important.  This  is  the  compound 
known  as  hydrogen  sulphide. 

Hydrogen  Sulphide,  or  sulphuretted  hydrogen,  occurs 
in  the  free  state  in  some  volcanic  gases  and  certain  mineral 
spring  waters  to  which  it  imparts  a  marked  odor. 
It  is  produced  by  the  decomposition  of  vegetable 
matters  containing  sulphur,  or  from  animal  albuminoids. 
It  can  be  formed  by  the  direct  union  of  the  two 
elements,  hydrogen  and  sulphur,  at  a  high  temperature, 
but  is  most  easily  made  by  a  method  quite  analogous  to 
that  by  which  hydrochloric  acid  is  made  from  a  chloride, 
that  is  by  the  decomposition  of  a  sulphide  by  means  of 
some  strong  acid.  We  use  for  this  purpose  an  artificial 
substance  known  as  the  sulphide  of  iron,  or  ferrous  sul- 
phide, the  preparation  of  which  was  illustrated  in  the 
experiment  wherein  iron  and  sulphur  were  melted  together 
and  strongly  heated. 

When  an  acid  is  poured  over  the  ferrous  sulphide  it  is 
decomposed  with  production  of  the  hydrogen  sulphide,  as 
illustrated  by  the  following  equation  : 

FeS  +  H2S04  =  H2S     -f  FeSO4 

Ferrous     I     Sulphuric    =r  Hydrogen     I   Ferrous 
sulphide    '  acid  sulphide      'sulphate. 

This  method  of  making  the  hydrogen  sulphide  is  shown 
in  the  following  experiment: 


166 


GENERAL  CHEMISTRY. 


Ex.  118.  Fit  a  bottle  with  a  funnel  tube  and  delivery  tube  as 
shown  in  the  figure  below.  Put  in  it  about  20  to  25  Gm.  of  ferrous  sul- 
phide in  small  lumps  and  pour  through  the  funnel  tube  enough  dilute 
sulphuric  or  hydrochloric  acid  to  cover  the  sulphide  and  the  end  of  the 
tube.  The  delivery  tube  leads  down  into  a  dry  bottle  covered  as  well  as 
possible  during  the  experiment  with  a  glass  plate.  The  gas  is  slightly 
heavier  than  air  and  can  therefore  be  collected  in  this  manner,  but  not 
perfectly.  After  adding  the  acid,  a  few  minutes  must  be  allowed  for  the 
expulsion  of  air  by  the  liberated  gas;  then  collect  several  bottles  of  the 
latter  for  experiment.  With  one  bottle  show  the  combustibility  of  the 
gas  by  burning  it  as  with  hydrogen ;  observe  the  odor  and  test  the  reaction 


FIG.  19. 


of  the  product  of  the  combustion  by  means  of  litmus  paper.  With  a 
second  bottle  of  the  gas  show  its  solubility  in  water.  To  this  end  cover 
the  bottle  with  a  glass  plate,  invert  it  with  the  mouth  under  water  and 
then  remove  the  plate.  After  a  time  water  ascends  into  the  bottle  as 
the  gas  is  absorbed.  All  this  work  must  be  done  in  a  fume  closet.  The 
gas  is  somewhat  poisonous. 

Ex.  119.  After  showing  the  properties  of  the  gas,  as  above,  re- 
place the  dry  bottles  by  one  containing  water.  As  the  gas  has  been 
found  to  be  to  some  degree  soluble.a  solution  known  as  hydrogen  sulphide 
water  is  thus  obtained.  While  this  solution  is  being  made,  prepare 
weak  aqueous  solutions  of  copper  sulphate,  lead  acetate,  and  zinc  ace- 


GENERAL  CHEMISTRY.  167 

tate  and  pour  them  into  small  flasks  or  beakers.  Twenty-five  Cc.  of 
each  solution  will  be  sufficient.  Remove  the  bottle  containing  the 
water  from  the  generator  and  in  its  place  put  the  flask  or 
beaker  with  the  copper  sulphate.  As  the  gas  bubbles  enter  this  solu- 
tion a  black  precipitate  forms.  In  time  this  precipitate  will  settle  to  the 
bottom  of  the  vessel.  Now  take  off  the  delivery  tube,  wash  and  replace  it 
and  let  the  gas  pass  next  into  the  solution  of  lead  acetate.  A  black  pre- 
cipitate forms  here  also.  Next,  after  washing  the  tube,  pass  the  gas  into 
the  zinc  solution.  A  precipitate  appears  here,  but  it  is  white  instead  of 
black.  After  a  few  minutes  add  to  the  precipitating  mixture  half  a 
dozen  drops  of  hydrochloric  acid  and  notice  that  the  precipitate  disap- 
pears or  is  dissolved  by  the  acid.  Then  from  a  test-tube  add  ammonia 
water,  a  drop  at  a  time,  and  observe  that  when  a  certain  amount  has 
been  added  the  precipitate  forms  again.  To  the  flasks  containing  the 
precipitates  from  the  lead  and  copper  solutions  add  a  little  hydrochloric 
acid  and  notice  that  the  precipitates  fail  to  disappear,  as  in  the  case  of 
that  from  the  zinc  solution.  After  these  tests  have  been  made,  allow 
the  remainder  of  the  gas  to  bubble  into  some  ammonia  water  contained 
in  a  small  flask.  This  yields  a  solution  of  ammonium  sulphide, 
(NH4)2S.  Make  all  these  experiments  in  the  fume  closet. 

Properties.  In  the  foregoing,  some  of  the  most  char- 
acteristic properties  of  the  hydrogen  sulphide  have  been 
shown.  It  dissolves  in  water  just  as  hydrochloric  acid 
does,  but  to  a  much  less  extent.  Most  of  the  gas  can  be 
easily  expelled  by  boiling.  One  volume  of  water  dissolves 
about  two  volumes  of  the  gas  at  the  ordinary  temperature. 
Under  pressure  the  gas  can  be  condensed  to  a  liquid  which 
boils  at  — 61°. 

In  the  combustion  of  the  gas  we  have  a  phenomenon 
reminding  of  the  behavior  of  hydrogen,  but  the  flame  is 
weaker  and  the  odor  of  the  product  of  combustion  charac- 
teristic. Sulphurous  oxide  is  formed  in  this  experiment,  as 
is  illustrated  by  the  equation: 

H2S+30  =  H20+S02. 

In  the  moist  condition  this  gas  gave  the  acid  test  with 
litmus  paper. 

The  most  important  behavior  of  the  gas  is  shown,  how- 
ever, by  the  precipitation  of  the  three  solutions.  These 
precipitates  are  called  sulphides  and  are  formed  by  a  double 
decomposition  between  the  hydrogen  sulphide  and  the 


168  GENERAL  CHEMISTRY. 

substances  in  solution.    The  following  equations  illustrate 
their  production: 

CuSO4-f   H2S   =  CuS    +  H2SO4 

Copper        (Hydrogen  __    Copper     i    Sulphuric 
sulphate      I     sulphide         sulphide   I         acid 

Pb(C2H3O2)2  +   H2S   =  PbS    -f2HC2H3O2 

Lead  (Hydrogen Lead        I  Acetic 

acetate  •    sulphide        sulphide    '  acid 

Zn(C2H3O2)2-f   H2S   =   ZnS  +  2HC2H3O2 

Zinc  _|_  Hydrogen —-      Zinc       •  Acetic 

acetate  '     sulphide        sulphide   I  acid. 

As  experiment  shows  and  as  the  equations  illustrate, 
the  solutions,  though  they  may  be  neutral  to  begin  with, 
must  become  acid  during  the  precipitation,  because  acids 
are  formed.  The  precipitates  are  therefore  insoluble  in 
these  acids.  That  the  sulphides  of  copper  and  lead  are  in- 
soluble in  hydrochloric  acid  also  was  shown  above,  but 
zinc  sulphide  was  found  to  be  soluble  in  hydrochloric  acid. 
This  behavior  is  very  suggestive;  we  are  able  by  it  to  dis- 
tinguish between  lead  and  zinc  solutions,  although  both 
may  be  colorless.  In  the  one  case  we  get  a  black  sulphide, 
insoluble  in  hydrochloric  acid,  while  in  the  other  we  ob- 
tain a  white  precipitate  soluble  in  hydrochloric  acid. 

To  more  fully  illustrate  this  important  matter  let  the 
student  carry  out  the  following  tests: 

Ex.  120.  Prepare  dilute  solutions  of  mercuric  chloride,  antimony 
chloride,  ferrous  sulphate,  zinc  sulphate,  calcium  chloride  and  sodium 
chloride.  Acidify  each  one  with  hydrochloric  acid  and  pass  in  hydrogen 
sulphide  gas  from  a  generator  as  before.  With  the  mercury  and  anti- 
mony solutions  we  obtain  precipitates,  in  the  first  case  black  and  in  the 
other  orange.  But  in  the  other  cases  no  precipitates  form,  even  after 
passing  the  gas  a  long  time.  While  the  gas  is  passing  add  to  each  solu- 
tion in  turn  some  ammonia  water,  and  note  the  result.  With  the  first 
there  is  apparently  no  change,  but  with  the  second  there  is.  The  orange 
yellow  precipitate  dissolves  to  form  a  dark  yellow  solution.  In  the  iron 
solution  we  have  a  black  precipitate  and  in  the  zinc  a  white,  while  in 
the  other  solutions  no  precipitate  appears.  Add  now  to  these  last  some 
solution  of  ammonium  carbonate.  In  the  calcium  solution  a  white  pre- 
cipitate is  formed  while  the  other  remains  clear. 

It  appears  from  the  above  that  the  mercury  compound 
yields  a  precipitate  in  the  presence  of  both  acid  and  alkali, 


GENERAL  CHEMISTRY.  169 

the  antimony  compound  'gives  a  precipitate  from  the  acid 
solution  only,  the  iron  and  zinc  compounds  from  the  alka- 
line solution  only,  while  the  calcium  and  sodium  com- 
pounds give  no  sulphide  precipitates.  The  iron  precipitate 
is  distinguished  from  the  zinc  precipitate  by  its  color,  while 
finally  the  calcium  and  sodium  compounds  which  give  no 
precipitates  with  the  gas  are  distinguished  from  each 
other  by  the  fact  that  one  yields  a  precipitate  with  am- 
monium carbonate,  while  the  other  does  not.  It  should  be 
added  here  that  zinc  may  be  precipitated  in  presence  of 
acetic  acid,  as  well  as  from  alkaline  solutions. 

These  experiments  are  of  fundamental  importance,  and 
the  student  will  learn  later  that  they  are  of  common  appli- 
cation in  the  branch  of  chemistry  known  as  qualitative 
analysis.  The  hydrogen  sulphide  is  one  of  our  most 
important  test  substances  and  by  its  aid  we  are  able  not 
only  to  recognize  bodies  in  solutions,  but  to  make  separa- 
tions1 of  bodies  into  groups,  and  thus  isolate  them  from  each 
other. 

We  have  remaining  our  solutions  of  hydrogen  sulphide 
in  water  and  in  ammonia  water.  With  these,  experiments 
may  be  made  to  show  that  they  behave  in  many  cases  as 
does  the  gas.  Let  the  student  determine  for  himself  their 
action  with  solutions.  Both  are  important  reagents  in  the 
laboratory. 

SULPHUR  AND  CHLORINE. 

Three  combinations  of  sulphur  with  chlorine  are 
known. 

Sulphur  Monochloride,  S2C12.  This  is  a  yellowish 
brown  liquid  easily  made  by  passing  dry  chlorine  over  sul- 
phur melted  in  a  retort.  The  liquid  has  a  specific  gravity 
of  1.705  and  boils  at  138°.  It  dissolves  sulphur  readily  and 
in  large  quantity.  The  solution  so  made  is  used  in  vul- 
canizing rubber. 

Sulphur  Dichloride,  SC12,  is  made  by  passing  dry 
chlorine  into  the  monochloride  at  a  temperature  kept  near 
the  zero  point.  It  is  a  dark  liquid  which  decomposes  if 


170  GENERAL  CHEMISTRY 

heated  to  about  20°  or  above,  yielding  free  chlorine  and  the 
monochloride. 

Sulphur  Tetrachloride,  SC14,  is  made  by  saturating 
the  dichloride  with  chlorine  at  a  temperature  of — 22°.  It 
is  a  light  yellow  liquid  which  decomposes  quickly  at  tem- 
peratures above  — 22° 

Sulphur  forms  also  compounds  with  iodine  and  bro- 
mine, but  they  are  not  important. 

SELENIUM. 

This  is  a  comparatively  rare  element  which  resembles 
sulphur  in  many  respects.  It  is  a  solid  substance  with  a 
dark  gray  color  and  melts  at  about  217°. 

The  compounds  of  selenium  resemble  those  of  sul- 
phur. The  best  known  are:  selenium  dioxide,  SeO2,  a  white 
crystalline  substance  which  dissolves  in  water  to  form 
selenous  acid,  H2SeO3;  selenic  acid,  H2SeO4,  made  by  oxi- 
dation of  a  selenite;  hydrogen  selenide,  H2Se,  a  gas  resem- 
bling H2S  and  made  by  decomposing  a  selenide  by  an  acid, 
and  which  precipitates  many  metals  as  does  the  sulphide. 

Selenium  burns  with  a  disagreeable  odor  described  as 
resembling  that  of  rotten  cabbage,  forming  an  oxide  of  un- 
known composition. 

TELLURIUM. 

Tellurium  is  another  rare  element  of  the  sulphur  group, 
found  usually  in  combination  as  a  telluride.  In  appear- 
ance it  resembles  the  metals,  but  behaves  chemically  as 
do  sulphur  and  selenium.  Its  specific  gravity  is  6.2  and 
its  melting  point  is  about  500°.  Its  most  important  com- 
pounds are:  hydrogen  telluride,  H2Te,  a  gas;  tellurium  di- 
oxide, TeO2,  a  white  solid  which  forms  tellurous  acid, 
H2TeO3,  with  water;  telluric  acid,  H2TeO4,  a  white  solid, 
soluble  in  water. 

The  three  elements  just  considered  constitute  an  inter- 
esting natural  group  in  which  the  properties  of  the  ele- 
ments themselves  and  of  their  compounds  are  functions  of 
the  corresponding  atomic  weights.  This  is  shown  in  the 
following  table : 


GENERAL  CHEMISTRY. 


171 


Sulphur. 

Selenium. 

Tellurium. 

Atomic  weight  
Specific  gravity. 

32.1 
2.05 

79.0 
4.6 

127.5 
6.2 

Melting  point   

114° 

217° 

500° 

Boiling  point  

440° 

665° 

Above  1,000° 

Hydrogen  compound 
ous  oxide 

H2S,  gas. 
SO  2  a  gas  be- 

H2Se, gas. 
SeO2,  a  solid 

H2Te,  gas. 
TeO2,a  crys- 

— ous  acid 

comes  liquid  at 
—8°,  soluble  in 
water. 

H2SO3      not 

readily  soluble 
in  water. 

talline        solid, 
slightly  soluble 
in  water. 

H2TeO3,       a 

—  ic  oxide 

stable    in    free 
state. 

SO3    volatile 

Not  known. 

solid. 
TeO8,    yellow 

—  ic  acid 

solid. 
H2SO4  liquid 

H2  SeO4, 

crystalline  solid. 
H2TeO4,  white 

volatile  acid  not 
decomposed  by 
HC1. 

heavy  colorless 
liquid,    decom- 
posed by  HC1. 

solid  mass,  with 
water  H2TeO4 
-f2H2O,     crys- 
talline solid. 

It  will  be  noticed  that  the  specific  gravities,  the  melt- 
ing points  and  the  boiling  points  of  the  elements  increase 
with  the  atomic  weights.  Also  that  the  compounds  become 
heavier  or  more  nearly  solids  in  the  same  order.  It  will  be 
pointed  out  later  that  similar  relations  exist  between  the 
members  of  other  groups  and  that  in  a  general  way  the 
properties  of  elements  are  closely  dependent  on  their 
atomic  weights. 


CHAPTER  VIII. 


SILICON  AND  BORON  AND  THEIR  COMPOUNDS. 

THESE  elements  occur  in  nature  combined  with  oxygen 
or  with  oxygen  and  metals. 

SILICON. 

This  is  one  of  the  very  abundant  elements  in  combina- 
tion and  is  found  as  the  oxide,  SiO2,  in  several  minerals  of 
which  the  most  common  is  quartz.  Other  substances,  flint, 
white  sand,  opal,  chalcedony  and  agate,  consist  essen- 
tially of  this  oxide.  In  many  silicates  the  element  is 
widely  distributed,  and  it  follows  oxygen  in  point  of  abun- 
dance in  the  earth's  crust. 

Preparation.  The  element  may  be  separated  by 
decomposing  one  of  its  compounds  by  potassium,  by  aid  of 
heat: 


It  is  left  after  this  reaction  as  an  amorphous  powder. 
If  this  is  melted  with  zinc  it  becomes  crystalline  as  the 
zinc  cools  and  may  be  secured  in  this  form  by  dissolving 
the  metal.  Its  specific  gravity  is  2.49  and  it  is  hard  enough 
to  scratch  glass.  As  the  methods  followed  in  the  isolation 
of  silicon  are  expensive  the  free  element  has  no  technical 
uses. 

SILICON  AND  OXYGEN. 

But  one  oxide  of  silicon  is  known  and  this  has  the  com- 
position SiO2,  and  is  called  silica.  It  occurs  in  very  pure 
form  in  varieties  of  quartz  and  tridymite,  both  characteris- 


GENERAL  CHEMISTRY.  173 

tic  crystalline  minerals.      In  the  opal   it  is  found  in  amor- 
phous condition. 

Preparation.  Pure  silica  may  be  easily  made  by  de- 
composing a  solution  of  sodium  silicate,  known  as  soluble 
glass,  by  means  of  hydrochloric  acid.  The  precipitate 
which  forms  is  thoroughly  washed  with  water,  dried  and 
ignited.  This  leaves  the  silica  in  a  fine  amorphous  condi- 
tion. 

Properties.  Silica  is  practically  insoluble  in  cold 
water  and  common  acids.  It  is  dissolved,  however,  by 
hydrofluoric  acid,  on  which  fact  the  etching  of  glass  de- 
pends. At  a  very  high  temperature  water  (superheated) 
dissolves  silica  to  a  slight  extent,  forming  silicic  acid.  In 
alkali  solutions,  especially  if  warm,  amorphous  silica  dis- 
solves readily  to  form  silicates.  The  crystalline  varieties 
of  silica  may  be  converted  into  the  same  silicates  by  fusion 
with  alkalies. 

SILICIC  ACIDS. 

Silicon  has  a  valency  of  four  and  the  acid  correspond- 
ing to  it  with  the  greatest  molecular  weight  has  the  for- 
mula H4O4Si,  and  is  known  as  orthosilicic  acid.  Its  com- 
position is  represented  in  this  manner: 


H— O\ 
H— O—  |  c. 
H-0/      Sl 
H— O7 


This  acid  exists  in  solution,  but  as  it  is  not  stable,  can- 
not be  obtained  in  the  free  state.  An  acid  derived  from 
this  is  known: 

H4SiO4— H2O  =  H2SiO3, 

H— Ox 

>Si  =  O. 
H— O7 

To  prepare  solutions  of  these  acids  a  weak  solution  of 
water-glass  is  decomposed  by  hydrochloric  acid,  leaving 


174  GENERAL  CHEMISTRY. 

orthosilicic  acid  dissolved.  The  mixture  is  thrown  on  a 
dialyzer,  floating  on  water,  and  allowed  to  remain  until  it 
is  free  from  hydrochloric  acid  and  chlorides.  These  sub- 
stances pass  through  the  membrane  bottom  of  the  dialyzer, 
but  the  colloidal  silicic  acid  cannot.  In  this  manner  it  is 
possible  to  prepare  a  weak, pure  solution  of  the  ortho  acid. 
This  may  be  concentrated  to  a  strength  of  about  14  per 
cent. 

When  evaporated  beyond  this,  water  is  lost  and  the 
gelatinous  acid,  H2SiO3,  is  formed.  This  in  turn  by  loss 
of  water  becomes  SiO2. 

H2Si03— H20: 


Silicates. 

Corresponding  to  the  silicic  acids  a  large  number  of 
bodies  called  silicates  are  known.  Some  of  these  can  be 
formed  artificially,  but  most  of  them  occur  in  nature  as 
mineral  species,  many  of  which  are  common  and  impor 
tant  bodies.  The  composition  of  most  of  these  minerals 
appears  quite  complex,  but  a  little  study  shows  their  rela- 
tion to  orthosilicic  acid.  For  instance,  the  mineral  serpen- 
tine may  be  represented  by  the  formula  Mg3Si2O7,  which 
corresponds  to  an  acid,  H6Si2O7.  Now,  this  in  turn  is 
related  to  the  ortho  acid,  as  illustrated: 

2H4SiO4  =  H6Si807+H80. 

The  common  mineral,  orthoclase,  is  called  a  trisilicate, 
and  is  represented  essentially  by  the  formula  AlKSi3O8, 
corresponding  to  H4Si3O8. 

This  is 

3H4Si04— 4H20  =  H4Si308. 

It  appears,  therefore,  that  these  silicates  may  be  looked 
upon  as  derived  from  condensed  silicic  acids,  formed  from 
orthosilicic  acid  by  loss  of  water.  It  will  be  pointed  out 
that  boric  acid  behaves  much  in  the  same  way. 

The  silicates  of  the  alkali  metals  are  soluble  in  water. 


GENERAL  CHEMISTRY.  175 

Potassium  silicate  and  sodium  silicate  are  called  soluble 
glass  or  water-glass. 

Ex.  121.  Take  about  5  Cc.  of  the  strong  solution  of  sodium  sili- 
cate, known  as  water-glass,  and  add  to  it,  a  little  at  a  time,  some  con- 
centrated hydrochloric  acid.  When  the  mixture  becomes  quite  strongly 
acid  a  gelatinous  mass  is  produced,  which  becomes  so  stiff  that  the 
test-tube  in  which  it  is  formed  may  be  inverted  without  spilling  it.  The 
thick,  colloidal  substance  is  impure  orthosilicic  acid  and  metasilicic 
acid. 

If,  before  adding  the  hydrochloric  acid,  the  water-glass 
is  largely  diluted  with  water  no  separation  of  the  colloidal 
substance  takes  place.  It  remains  in  solution  and  can  be 
partially  purified  by  dialysis,  as  explained  above. 

Silicic  acid  forms  insoluble  salts  with  many  basic  bod- 
ies, and  some  of  these  can  be  made  by  precipitation,  as 
shown  below  : 

Ex.  122.  Dilute  the  common  water-glass  with  about  20  parts  of 
water.  Take  small  portions  of  this  diluted  liquid  in  test-tubes  and  add 
to  them  solutions  of  calcium  chloride,  copper  sulphate,  lead  nitrate  and 
cobalt  nitrate.  Precipitates  are  formed  which  are  silicates  of  the 
metals  in  these  salts. 

The  soluble  glass  has  approximately  the  composition 
Na2SiO3  and  the  insoluble  silicates  may  be  made  from  it 
by  double  decomposition,  as  : 


Na2SiO3+CaCl2=:CaSiO3+2NaCl 

Sodium        (Calcium Calcium     _|_  Sodium 

silicate        'chloride  —    silicate        '     chloride. 


A  soluble  salt  is  left  in  the  liquid.  Advantage  is  taken 
of  this  behavior  of  the  soluble  silicate,  or  water-glass,  in 
making  certain  kinds  of  cement  and  artificial  stone. 

As  mentioned,  the  alkali  silicates  are  soluble;  the 
others  are  insoluble  in  water  and  many  of  them  cannot  be 
decomposed  by  acids.  Common  glass  is  an  artificial  mix- 
ture of  silicates  made  by  fusing  quartz  sand,  silica,  with 
basic  substances.  For  example,  common  window  glass  is 
made  by  fusing  a  mixture,  in  certain  proportions,  of  sand, 
lime  or  pure  limestone  and  dry  sodium  carbonate. 

Roughly    speaking,   we   distinguish   four  varieties    of 


176  GENERAL  CHEMISTRY. 

glass,  viz.  :    crown  or  window  glass,  Bohemian  glass,  flint 
glass  and  common  bottle  glass. 

Crown  Glass  is  essentially  a  mixture  of  calcium  and 
sodium  silicates.  In  some  kinds  a  little  alumina  is  present. 
It  is  made  by  melting  at  a  very  high  heat  a  mixture  of 
white  sand,  lime  or  limestone,  and  soda  ash  or  dry  sodium 
sulphate.  Common  window  and  plate  glass  and  much 
hollow  ware  are  included  under  crown  glass. 

Bohemian  Glass  consists  essentially  of  the  silicates  of 
potassium  and  calcium.  It  is  made  of  carefully  selected 
materials,  usually  quartz  sand,  pure  refined  potassium  car- 
bonate and  chalk,  or  well  burned  lime,  as  free  as  possible 
from  magnesia.  This  glass  can  be  fused  only  at  a  high 
temperature,  and  softens  only  with  difficulty  when  heated. 
It  is,  therefore,  employed  in  making  much  chemical  glass- 
ware. Sometimes  a  little  sodium  carbonate  is  used  with  the 
potassium  carbonate  to  make  it  more  readily  workable. 

Flint  Glass  is  essentially  a  lead  potassium  silicate 
and  is  made  by  melting  a  combination  of  sand,  red  lead 
and  dry  potassium  carbonate.  This  glass  can  be  melted 
and  cast  or  otherwise  worked  with  comparative  ease,  and 
is  therefore  employed  in  making  tableware  and  large  arti- 
cles of  ornamentation.  The  ready  fusibility  depends  on  the 
presence  of  lead  silicate.  This  glass  cannot  be  used  for 
chemical  ware. 

Common  Green  Bottle  Glass  resembles  crown  glass, 
but  is  made  of  impure  materials.  It  usually  contains  con- 
siderable quantities  of  iron  and  aluminum  silicates.  The 
green  color  is  due  to  the  ferrous  salt. 

The  chemical  composition  of  several  kinds  of  glass  as 
found  by  analysis  is  given  in  the  following  table,  the  re- 
sults being  expressed  in  the  usual  manner.  It  must  be 
remembered,  however,  that  certain  special  kinds  of  glass 
contain  still  other  substances.  A  part  of  the  silicic  acid 
may  be  replaced  by  boric  acid  and  for  some  purposes  the 
oxide  of  lead  may  be  partly  replaced  by  oxide  of  thallium. 
The  relation  of  glass  to  pottery  will  be  shown  later. 


GENERAL  CHEMISTRY. 


177 


KIND    OF    GLASS. 

Si02 

Na20 

K20 

CaO 

MgO 

PbO 

Fe203 

A1203 

BOHEMIAN. 

Combustion  tubing  

74.19 
76.41 

1.87 
1.38 

13.13 
10.96 

9.39 
9.71 

0.36 

0.49 
0  89 

Optical  glass 

75  81 

2  00 

15  03 

12  13 

0  32 

1  02 

Mirror  plate  

75.81 

4.84 

1L39 

7.38 

0.10 

1,01 

FLINT  GLASS. 

German  

75.24 

12.51 

1.48 

10.48 

English 

51  40 

9  40 

37  40 

2  00 

Optical  

44.30 

11.75 

43.05 

0.50 

0,12 

CROWN  GLASS. 

German  window  

71.56 

12.97 

13.27 

1. 

29 

English  window. 

70  71 

13  25 

13  38 

1. 

02 

French  plate  

73.00 

11.50 

15.50 

English  plate       . 

77  90 

12  53 

1.72 

4.85 

3. 

59 

German  plate  

78.75 

13.00 

6.50 

1. 

75 

Coloring  Glass.  Certain  metallic  oxides  may  be 
melted  with  the  glass  mixture  and  so  impart  to  the  finished 
glass  some  desired  shade.  The  red  oxide  of  copper  is  used 
in  making  ruby  glass,  while  the  black  oxide  of  the  same 
metal  gives  a  green  color;  ferrous  oxide  yields  a  green 
glass  and  ferric  oxide  a  yellowish  brown;  the  black  oxide 
of  manganese  is  used  in  giving  a  pink  to  purple,  oxide  of 
cobalt  a  deep  blue,  oxide  of  uranium  a  beautiful  canary 
yellow.  Various  shades  may  be  made  by  properly  com- 
bining some  of  these  oxides,  and  it  is  also  possible  by  the 
proper  combination  to  secure  from  impure  materials  an 
almost  colorless  glass.  Black  oxide  of  manganese  is  com- 
monly employed  to  correct  the  objectionable  color  due  to 
presence  of  iron.  This  it  does  by  oxidizing  the  iron  to  the 
ferric  condition,  the  yellow  tint  of  which  is  complementary 
to  the  purple  of  the  manganese  compound. 

SILICON  AND  HYDROGEN. 

One  compound  of  these  two  substances  is  known,  hav- 
ing the  composition  SiH4.  It  is  a  gaseous  body,  made  by 
the  action  of  acids  on  magnesium  silicide  and  is  an  inter- 
esting compound  from  a  theoretical  standpoint,  but  has  no 
technical  applications. 


178  GENERAL  CHEMISTRY. 

SILICON  AND  THE  HALOGENS. 

Silicon  exists  in  combination  with  fluorine,  chlorine, 
bromine  and  iodine.  Of  these  the  tetrafluoride,  SiF4,  is 
the  most  important.  The  formation  of  this  in  the  etching 
of  glass  has  been  referred  to  already.  It  is  produced  by 
the  action  of  hydrofluoric  acid  on  silica: 

4HF+Si02=SiF4+2H20. 

It  is  a  gaseous  substance  which  is  decomposed  by  con- 
tact with  water. 

The  compounds  SiHCl3  and  SiCl4  are  known.  They 
are  volatile  liquids  which  decompose  when  mixed  with 
water. 

Fluosilicic  Acid.  When  silicon  tetrafluoride  is  passed 
into  water  it  decomposes  in  this  way: 

2SiF4+3H2O  =  H2SiF6-f2HF+H2SiO3. 

The  body,  H2SiF6,  is  known  as  fluosilicic  acid.  It  is 
stable  only  in  solution,  and  in  this  form  is  sometimes  used 
as  a  reagent. 

BORON. 

This  is  an  element  which  is  found  in  a  few  natural  sub- 
stances, of  which  borax,  boric  acid  and  calcium  borate  are 
the  most  important.  The  element  may  be  liberated  by  the 
decomposition  of  some  of  its  compounds,  but  it  is  not  im- 
portant in  the  free  state.  The  specific  gravity  of  crystal- 
lized boron  is  2.68. 

BORON  AND  OXYGEN. 

One  compound  is  known  having  the  composition  B2O3. 
It  is  a  glass  like  body,  soluble  in  water,  made  best  by 
strongly  heating  boric  acid. 

BORIC  ACID. 

This  is  a  combination  of  boron  with  hydrogen  and  oxy- 
gen, having  the  composition  H3BO3.  It  is  found  in  na- 


GENERAL  CHEMISTRY.  179 

ture  in  small  amount,  and  especially  in  the  vapor  from  vol- 
canic fissures  existing  in  certain  parts  of  Tuscany.  The 
water  condensed  from  this  vapor  is  collected  in  small 
lagoons,  and  kept  boiling  by  the  action  of  the  hot  vapor 
itself.  In  this  way  a  rapid  concentration  is  effected. 
Much  of  the  boric  acid  of  commerce  comes  from  this 
source.  It  can  be  made  from  borax,  however,  and  this 
will  be  illustrated  here. 

Ex.  123.  By  the  aid  of  heat  dissolve  about  30  Gm.  of  powdered 
borax  in  about  120  Cc.  of  water.  Add  to  tne  hot  solution  enough  strong 
hydrochloric  acid  to  make  the  liquid  strongly  acid  to  litmus  paper.  Stir 
well  while  adding  the  acid.  Then  allow  the  mixture  to  cool  thoroughly. 
Thin  crystalline  plates  of  boric  acid  separate.  Remove  the  supernatant 
liquid  by  filtration,  take  up  the  boric  acid  with  hot  water  and  purify  it 
by  recrystallization. 

Ex.  124.  Boric  acid  is  much  more  soluble  in  hot  water  than  in 
cold.  It  is  also  readily  soluble  in  alcohol.  Prove  this  by  dissolving  the 
product  of  the  last  experiment  in  some  alcohol.  Pour  some  of  the  solu- 
tion so  obtained  over  a  little  asbestos  in  a  porcelain  dish.  When  the 
asbestos  is  thoroughly  moistened  take  it  up  with  clean  forceps  and  hold  it 
in  the  flame  of  the  Bunsen  burner  to  ignite  the  alcohol.  The  flame 
produced  is  colored  an  intense  green  by  the  hot  vaporized  boric  acid. 
This  is  a  characteristic  reaction  and  is  employed  for  the  recognition  of 
boron  compounds. 

Considerable  quantities  of  boric  acid  and  borax  are  now 
made  from  borocalcite,  CaB4O7.4H2O,  which  occurs 
abundantly  in  California. 

Borax.  This  substance  is  a  salt  containing  boric  acid 
and  sodium.  Its  chemical  composition  is  shown  by  the 
formula  Na2B4O7-f-10H2O.  The  water  here  represented 
is  known  as  water  of  crystallization,  and  can  be  separated 
by  heat,  leaving  what  is  known  as  borax  glass  or  anhydrous 
borax.  The  preparation  of  this  can  be  shown  by  a  simple 
test. 

Ex.  125.  Bend  the  end  of  a  piece  of  platinum  wire  so  as  to  form  a 
loop  two  or  three  millimeters  across.  Heat  this  in  the  flame  of  the  Bun- 
sen  burner  and  then  dip  it  while  hot  in  some  powdered  borax.  Heat  again 
and  repeat  the  operation  until  the  loop  is  well  covered.  Then  hold  this 
in  the  flame  several  minutes.  The  mass  swells  and  gives  off  steam,  but 
finally  fuses  together  and  forms  a  colorless,  clear,  glass-like  globule, 
filling  the  loop,  and  called  the  borax  bea<i.  The  composition  of  this  is 


180  GENERAL  CHEMISTRY. 

Na2B4O7.  This  "bead  has  a  very  important  application  in  analytical 
chemistry  which  is  illustrated  in  this  manner:  While  it  is  still  warm  bring 
it  in  contact  with  a  minute  particle  of  cobalt  oxide  or  other  compound 
of  this  metal.  The  substance  sticks  to  the  hot  borax,  and  when  the 
bead  is  again  fused  in  the  flame  dissolves  in  it  completely,  imparting  a 
blue  color,  best  seen  when  the  bead  cools.  Many  other  metals  are  dis- 
solved from  their  combinations  by  the  borax  in  the  same  manner, 
imparting  some  color  which  is  characteristic  of  the  metal.  It  will  be 
shown  later  that  nickel  gives  a  yellowish  brown  color,  manganese  a  pink, 
copper  a  bluish  green  color,  and  other  metals  other  shades. 

The  amount  of  copper,  cobalt,  nickel,  manganese  or 
other  body  necessary  to  impart  color  to  the  borax  bead  is 
very  small,  and  hence  the  test  is  a  very  delicate  one  for  the 
recognition  of  traces  of  these  substances.  The  solution 
of  these  different  substances  in  the  fused  borax  depends 
on  the  formation  of  new  compounds  which  are  known  as 
double  borates.  The  boric  acid  which  enters  into  the  com- 
position of  borax  is  capable  of  holding  or  uniting  with  a 
much  greater  proportion  of  the  basic  substance  than  is 
already  present.  Hence  we  have  sodium-cobalt  borate, 
sodium-nickel  borate  and  so  on. 

The  explanation  of  this  interesting  behavior  will  appear 
if  we  look  into  the  nature  of  boric  acid  itself.  The  struc- 
ture of  boric  acid  is  represented  in  this  way: 

H— O^ 
H— O— B 
H— O/ 

This  is  called  orthoboricacid,  as  thecompound  H4SiO4 
is  called  orthosilicic  acid,  and  the  various  borates  are 
formed  by  replacing  the  hydrogen  by  metallic  atoms.  The 
orthoborates  are  represented  by  the  composition  M3O3B, 
where  M  stands  for  an  atom  with  unit  valence. 

These  orthoborates,  however,  are  not  common  sub- 
stances, but  the  so-called  metaborates  are  more  stable 
and  more  common.  Metaboric  acid  is  derived  from  the 
ortho  acid  by  loss  of  water: 

H3BO3— H2O  =  HBO2. 
Finally  we  have   another   acid  still    more  condensed, 


GENERAL  CHEMISTRY.  181 

known  as  pyroboric  acid,  which  may  be  made  by  heating 
boric  acid  to  140°  for  some   hours: 

4H3B03=H2B4074-5H20. 

This  new  acid  is  a  glassy,  brittle  mass.  Its  structure 
may  be  represented  in  this  way: 

H—  .0—  B/     \B  —  O  —  B/    \B—  O—  H. 

XO/  XO/ 

The  sodium  salt,  or  borax,  contains  sodium  instead  of 
the  hydrogen,  and  we  can  represent  the  probable  struc- 
ture as  above,  or  as  below,  which  amounts  to  the  same: 

Na—  O—  B 


Na-Q-B 

Now,  when  copper  oxide,  cobalt  oxide  or  some  other 
metallic  compound  is  fused  with  borax  a  union  takes  place 
so  that  bodies  of  this  type,  possibly,  are  formed: 

Na—  O— 


Na  —  O—  -W 

That  is,  the  dissolved  metal  takes  the  place  of  hydrogen 
in  a  derived  acid.  Because  of  this  property  of  dissolving 
oxides  borax  is  used  by  the  blacksmith  in  welding,  by  the 
jeweler  in  soldering  and  6y  the  assayer  as  a  flux  in  various 
metallurgical  operations. 


182  GENERAL  CHEMISTRY. 

Other  Boron  Compounds.  Boron  is  found  in  com- 
bination with  the  halogen  elements,  in  bodies  of  the  type 
BC13.  It  forms  a  singular  compound  with  nitrogen,  which 
has  the  structure,  BN.  It  is  a  white  solid  decomposed  by 
water  in  this  way: 

BN+3H2O  =  H3BO3-fNH3. 

Detection  of  Boron.  This  element  is  usually  recog- 
nized in  its  compounds  by  the  flame  test,  as  explained 
above.  The  compound  is  mixed  with  a  little  sulphuric 
acid,  to  liberate  boric  acid,  and  alcohol.  When  the  latter 
is  ignited,  as  explained,  the  flame  shows  a  green  color  in 
presence  of  boric  acid.  Boric  acid  is  also  recognized  by 
the  brown  color  it  imparts  to  paper  moistened  with  solu- 
tion of  turmeric.  This  color  is  not  discharged  by  hydro- 
chloric acid. 


CHAPTER  IX. 


PHOSPHORUS     AND     ARSENIC     AND     THEIR 
COMPOUNDS. 

PHOSPHORUS. 

Occurrence.  This  element  is  never  found  free  in  nature, 
but  occurs  always  combined  with  oxygen  and  metals. 
It  is  most  abundant  in  calcium  phosphate,  Ca3(PO4)2, 
which  is  found  in  bones  and  in  several  minerals,  and 
also  in  amorphous  deposits.  There  are  also  a  number 
of  other  minerals  which  contain  phosphorus  ;  it  is  found 
in  several  of  the  body  tissues  and  is  excreted  as  a  phos- 
phate in  the  urine. 

History.  Phosphorus  was  first  obtained  by  Brand,  an 
alchemist  of  Hamburg,  about  the  middle  of  the  seventeenth 
century.  It  was  produced  from  urine  in  an  attempt  to  dis- 
cover a  process  for  the  conversion  of  base  metals  into  gold. 
The  general  process  was  somewhat  improved  by  others,  but 
for  years  the  price  of  the  element  remained  very  high.  In 
1771  Scheele  showed  how  it  could  be  made  from  bone  ash, 
in  which  it  was  discovered  a  few  years  before.  Phosphorus 
is  now  used  for  several  purposes,  so  that  its  manufacture 
has  become  an  important  industry. 

Preparation.  On  burning  bones  a  white  residue, 
known  as  bone  ash,  and  consisting  mainly  of  calcium  phos- 
phate, is  left.  From  this  the  phosphorus  of  commerce 
has  been  usually  obtained  by  the  following  general  method. 
The  bone  ash  is  mixed  with  crude  sulphuric  acid,  some- 
what diluted,  which  brings  about  a  decomposition  illustrated 
by  the  following  equation  : 

Ca3(P04)2+2H2S04  =  CaH4(P04)2+2CaS04. 


184  GENERAL  CHEMISTRY. 

The  insoluble  calcium  phosphate  is  converted  into  a 
soluble  compound  containing  less  calcium,  and  known  as 
calcium  hydrogen  phosphate.  The  solution  of  this, 
CaH4(PO4)2,  is  separated  from  the  insoluble  calcium 
sulphate,  CaSO4,  and  is  evaporated  to  dryness  and  then 
strongly  heated.  By  this  treatment  water  is  driven  off  and 
a  compound  known  as  calcium  metaphosphate  is  left : 

CaH4(P04)2=Ca(P03)2+2H20. 

The  calcium  metaphosphate  is  then  mixed  with  char- 
coal and  distilled  from  clay  retorts.  Free  phospnorus 
passes  over  in  the  form  of  vapor,  which  is  condensed  under 
water  to  a  liquid  which  is  afterward  run  into  molds  and 
thus  formed  into  sticks  as  seen  in  commerce.  The  chem- 
ical changes  occurring  in  the  distillation  of  the  mixture 
from  the  clay  retorts  have  been  found  to  take  place  in  a 
manner  illustrated  by  this  equation  : 

3Ca(TO3)2+10C  =  P4H-Ca3(PO4)2-flOCO. 

The  action  of  the  charcoal  here  is  that  of  a  reducing 
agent,  to  remove  oxygen. 

At  the  present  time  phosphorus  is  largely  produced  in 
the  electric  furnace  by  decomposing  a  mixture  of  a  phos- 
phate and  powdered  coke  or  charcoal.  The  intense  heat 
of  the  electric  arc  makes  this  method  of  liberating  phos- 
phorus a  very  simple  one.  From  the  electric  furnace  the 
phosphorus  distills  and  is  condensed  under  water,  as  in  the 
older  process. 

Properties  of  Phosphorus.  In  perfectly  pure  con- 
dition phosphorus  is  a  white  wax-like  solid.  The  commer- 
cial product  is  always  yellowish.  It  may  be  cut  readily 
with  a  knife.  It  melts  at  a  temperature  of  44. 2°  C.  and  has  a 
specific  gravity  of  about  1.83.  It  boils  at  about  290°. 
Whenheated  to  a  certain  temperature  in  an  atmosphere  free 
from  oxygen  it  forms  red  phosphorus,  which  will  be  referred 
to  below.  Other  properties  have  been  illustrated  already  in 
the  experiments  with  oxygen,  and  in  the  analysis  of  air. 
Phosphorus  was  found  to  combine  very  readily  with  oxygen. 


GENERAL  CHEMISTRY.  185 

It  combines  directly  with  other  elements,  as  illustrated  in 
the  next  experiments. 

Ex.  126.  Place  a  very  small  piece  of  dry  phosphorus  in  a  porcelain 
crucible  or  evaporating  disband  add  a  few  small  crystals  of  iodine  in 
such  a  manner  as  to  bring  the  two  substances  into  actual  contact.  Very 
soon  a  combination  takes  place  accompanied  by  light  and  heat. 

Ex.  127.  «%In  a  dry  flask  or  bottle  with  a  wide  neck  warm  a  few 
drops  of  bromine  until  the  vapor  fills  the  vessel.  Put  a  small  piece  of 
dry  phosphorus  in  a  deflagrating  spoon  and  immerse  this  in  the  vapor. 
The  phosphorus  combines  quickly  with  the  bromine. 

In  the  first  of  these  experiments  iodide  of  phosphorus 
is  formed,  and  in  the  second  bromide  of  phosphorus.  Com- 
pounds of  chlorine  are  known,  made  by  direct  union. 

Solubility  of  Phosphorus.  Common  phosphorus  is 
not  soluble  in  water,  but  is  dissolved  by  several  other 
liquids.  One  of  the  best  solvents  is  carbon  disulphide, 
CS2.  The  action  is  shown  in  the  following  experiment: 

Ex.  128.  Pour  about  5  cubic  centimeters  of  carbon  disulphide  into 
a  test-tube,  and  add  a  very  small  piece  of  phosphorus,  not  larger  than  a 
small  pea.  Cork  the  tube  and  set  it  aside  a  few  minutes,  with  occa- 
sional shaking.  The  phosphorus  soon  dissolves.  Put  a  sheet  of  filter 
paper  on  a  brick  or  flat  stone,  and  pour  the  solution  from  the  test-tube 
over  the  paper.  Leave  none  of  it  in  the  tube.  Allow  the  paper  to  stand 
for  the  spontaneous  evaporation  of  the  disulphide.  This  is  complete  in 
a  few  minutes,  and  then  the  phosphorus,  spread  over  the  whole  paper, 
bursts  into  flame,  combining  suddenly  with  the  oxygen  of  the  air.  Be- 
fore putting  the  test-tube  away  heat  it  to  oxidize  any  trace  of  phosphorus 
left  in  it,  to  avoid  accidents. 

Common  phosphorus  is  violently  poisonous  when  taken 
into  the  stomach,  and  many  fatal  cases  are.on  record.  This 
property  is  probably  due  to  its  great  affinity  for  oxygen,  as 
the  oxidized  compounds  of  the  element  are  nonpoisonous. 

Uses  of  Phosphorus.  It  is  largely  employed  in  the 
manufacture  of  matches,  the  heads  of  which,  in  some  cases, 
consist  of  phosphorus  and  gum  arabic.  Because  of  the 
danger  to  workmen  in  factories  where  matches  are  made, 
the  use  of  common  phosphorus  has  been  in  many  places 
abolished.  The  red  variety,  to  be  described,  is  often  used 


186  GENERAL  CHEMISTRY. 

instead.     Phosphorus  is  also  largely  used  in  the  manufac- 
ture of  several  of  its  important  compounds. 

Amorphous  Phosphorus. 

A  peculiar  modification  of  phosphorus  exists,  known  as 
red  or  amorphous  phosphorus.  This  substance  has  few  cf 
the  marked  properties  found  in  the  ordinary  variety.  It  is 
not  poisonous  and  is  insoluble  in  carbon  disulphide. 
When  heated  to  a  temperature  of  about  240°  C.  the  com- 
mon white  phosphorus  is  changed  to  the  red  variety,  and 
the  latter  at  a  temperature  of  about  300°  is  changed  into 
the  white  form  again.  This  can  be  illustrated  by  an  ex- 
periment. 

Ex.  129.  Heat  a  little  red  phosphorus  in  a  narrow  glass  tube  sealed 
at  the  lower  end.  At  the  proper  temperature  a  light  colored  product  is 
formed,  which  vaporizes  and  collects  on  the  cooler  part  of  the  tube 
above  the  flame.  That  this  is  the  ordinary  phosphorus  can  be  shown  by 
breaking  off  the  sealed  end  of  the  tube  and  heating  again.  Air  now 
having  access,  combustion  of  the  phosphorus  soon  follows. 

Red  phosphorus  may  be  also  made  by  heating  the  com- 
mon variety  in  closed  vessels  to  a  temperature  of  10°  above 
its  boiling  point.  The  change  then  is  very  rapid.  This 
form  of  phosphorus  may  be  used  with  perfect  safety  in  the 
match  industry  and  is  largely  prepared  for  that  purpose. 
It  has  a  specific  gravity  of  2.1  and  is  as  hard  as  marble. 
In  commerce  it  appears  usually  in  the  powdered  form. 

PHOSPHORUS  AND  HYDROGEN. 

Three  compounds  of  these  elements  are  known,  called 
hydrogen  phosphides: 

PH3,    a  gas. 
P2H4,  a  liquid. 
P4H2,  a  solid. 

The  first  one  is  very  easily  made,  and  is  usually  accom- 
panied by  a  small  amount  of  the  second  or  liquid  product. 
The  latter  is  spontaneously  inflammable,  and  ignites  the 
other  on  coming  in  contact  with  the  air.  The  preparation 


GENERAL  CHEMISTRY. 


187 


of  the  gaseous  product  is  shown  in  the  experiment  now  to 
be  given: 

Ex.  130.  Arrange  the  apparatus  as  shown  in  the  figure.  The  flask 
has  a  capacity  of  about  250  to  300  Cc.  Two-thirds  fill  it  with  a  strong 
solution  of  potassium  hydroxide,  and  throw  in  a  small  piece  of  phos- 
phorus not  larger  than  a  pea.  Add  a  few  drops  of  ether,  insert  the  cork 
with  the  bent  delivery  tube,  and  apply  heat  gradually.  The  ether  on 
evaporating  serves  to  drive  the  air  out  from  the  upper  part  of  the  flask. 
The  delivery  tube  is  bent  so  that  its  lower  end,  which  is  turned  upward, 
dips  beneath  the  surface  of  water  in  the  basin.  When  the  liquid  be- 
comes warm  a  combination  between  the  alkali  and  the  phosphorus  takes 
place,  by  which  the  gaseous  hydrogen  phosphide  is  liberated.  This, 
\\ithatrace  of  the  liquid  product,  escapes  through  the  delivery  tube, 
and  as  soon  as  it  bubbles  through  the  water  into  the  air,  ignites  spon- 
taneously. In  the  combustion  phosphoric  oxide  is  formed,  which  takes 
the  shape  of  a  ring,  expanding  as  it  ascends  through  the  air.  Do  not  re- 


FIG.  20. 

move  the  lamp  or  the  tube  from  the  water  until  the  reaction  is  termi- 
nated. Then  lift  the  tube  from  the  water  first,  remove  the  lamp,  and 
allow  the  apparatus  to  cool.  The  gas  is  poisonous,  therefore  the  fittings 
must  be  tight. 

The  phosphorus  and  alkali   react  on  each  other  in  a 
manner  shown  by  the  following  equation  : 

P4+3H2O-f3KOH  =  3KH2PO2-f-PH3. 

Potassium 
hypophosphite 


188  GENERAL  CHEMISTRY. 

When  the  gas  burns  in  the  air  we  have: 
2PH3+402=P205+3H20. 

The  gas  PH3  is  commonly  called  phosphine.  It  forms 
several  combinations  resembling  those  formed  by  ammonia, 
and  these  are  known  as  phosphonium  compounds;  phos- 
phonium  bromide  is  formed  by  the  direct  combination  of 
the  gases  PH3  and  HBr, 

PH3+HBr=PH4Br. 

The  spontaneously  inflammable  gas  is  made  also  by 
decomposing  calcium  phosphide  with  water.  In  the  reac- 
tion some  liquid  hydrogen  phosphide  is  formed,  and  this 
may  be  separated  by  passing  the  gaseous  mixture  through 
tubes  immersed  in  a  freezing  bath.  A  colorless  liquid  con- 
denses and  this  burns  spontaneously  on  exposure  to  the 
air.  This  liquid,  P2H4,  is  not  stable  but  decomposes, 
yielding  P4H2,  a  yellow  solid,  and  PH3. 

PHOSPHORUS  AND  OXYGEN. 

Phosphorus  forms  four  or  five  oxides,  but  only  two  of 
them  are  important.  They  are  the  trioxide,  P2O3,  and  the 
pentoxide  P2O5.  Corresponding  to  these  oxides  are  two 
acids,  phosphorous  acid,  H3PO3,  and  phosphoric  acid, 
H3PO4.  A  third  acid  known  as  hypophosphorous  acid 
has  the  composition  H3PO2. 

Hypophosphorous  Acid.  A  salt  of  this  acid  was  pro- 
duced in  the  last  experiment.  In  the  flask,  after  the  com- 
pletion of  the  reaction,  we  have  the  solution  of  a  substance 
known  as  potassium  hypophosphite.  Other  bodies  called 
hypophosphites  are  made  on  the  large  scale  by  replacing 
the  potassium  hydroxide  by  different  alkali  substances. 
Some  of  these  products  are  very  important  and  are  used  in 
medicine.  The  free  acid  is  made  by  decomposing  the 
hypophosphite  of  barium  by  the  proper  amount  of  dilute 
sulphuric  acid: 


GENERAL  CHEMISTRY.  189 

A  concentrated  thick,  syrupy  liquid  is  obtained  by  evap- 
orating the  filtrate  from  the  barium  sulphate. 

The  acids  and  the  salts  are  all  strong  reducing  agents, 
because  of  the  tendency  of  the  phosphorus  to  combine 
with  more  oxygen.  A  solution  of  sodium  hypophosphite 
is  often  used  in  analytical  chemistry  for  the  precipitation 
of  certain  metals.  In  the  salts  of  hypophosphorous  acid 
but  one  atom  of  hydrogen  of  the  acid  may  be  replaced; 
they  are  all,  therefore,  of  the  type  H2NaPO3. 

Phosphorus  Trioxide.  This  is  a  white  powder, 
made  by  the  incomplete  combustion  of  phosphorus  in  a 
stream  of  air  or  oxygen  diluted  with  carbon  dioxide.  It 
dissolves  in  water,  forming  the  acid  known  as  phosphor- 
ous acid,  H3PO3: 


Neither  the  oxide  nor  the  acid  is  important.  The  salts 
of  the  acid  are  known  as  phosphites.  In  most  of  them  but 
two  of  the  hydrogen  atoms  of  the  acid  have  been  replaced. 
Phosphorous  acid  may  be  made  in  purest  form  by  decom- 
posing the  chloride,  PC13,  with  water. 

PCl3-}-3H2O  =  H3PO3+3HCl. 

Phosphorus  Pentoxide.  This  is  the  common  oxide 
produced  when  phosphorus  burns  in  the  air  or  in  a  suffi- 
cient supply  of  oxygen.  It  is  a  soft,  white  substance, 
characterized  by  its  powerful  affinity  for  water,  in  which  it 
dissolves,  forming  phosphoric  acid.  It  is  produced  in 
large  quantities  for  use  in  the  drying  of  gases,  for  several 
chemical  reactions  in  organic  chemistry,  and  also  as  a  step 
in  the  manufacture  of  phosphoric  acid  from  phosphorus. 

Phosphoric  Acids.  When  the  pentoxide  dissolves  in 
cold  water  this  reaction  follows: 


P2O5-fH2O  = 
The  substance  produced  here  is  known  as   metaphos- 


190  GENERAL  CHEMISTRY. 

phoric  acid.  It  is  not  stable  permanently  in  solution;  on 
standing  more  water  is  taken  up,  and  the  orthophosphoric 
acid,  H3PO4,  is  formed.  This  change  follows  very  quickly 
by  boiling  the  solution  of  the  metaphosphoric  acid. 

Orthophosphoric  acid  may  be  made  as  just  described, 
or  it  may  be  produced  by  the  oxidation  of  phophorus  by 
.means  of  nitric  acid.  It  is  also  largely  produced  from  pure 
bone  ash,  mainly  calcium  phosphate,  by  separation  of  the 
bases.  The  pure  acid  appears  in  commerce  as  a  thick, 
syrupy  liquid,  which  sometimes  deposits  crystals  on  stand- 
ing. This  is  the  common  phosphoric  acid  of  the  pharma- 
copoeia and  is  much  used  in  medicine.  It  mixes  with  water 
in  all  proportions,  and  is  stable  at  ordinary  temperatures. 
At  a  temperature  a  little  above  200°  the  pure  ortho  acid 
loses  water  and  yields  an  acid  known  as  pyrophosphoric 
acid,  H4P2O7. 

2H3PO4— H  ,O  =  H4P8  O7. 

Pyrophosphoric  acid  is  a  crystalline  body  easily  solu- 
ble in  water.  All  of  the  hydrogen  may  be  replaced  to 
form  salts,  some  of  which  are  quite  important.  Sodium 
pyrophosphate,  Na4P2O7,  is  most  easily  made  by  heating 
the  common  phosphate,  HNa2PO4,  to  redness.  A  mole- 
cule of  water  is  driven  off  and  the  pyrophosphate  results. 
Pyrophosphates  yield  with  silver  nitrate  a  white  precipi- 
tate, Ag4P2O7,  while  orthophosphates  yield  a  yellow  pre- 
cipitate of  Ag3PO4.  The  common  pyrophosphates  are 
stable  in  cold  solution;  on  long  boiling  they  revert  to  or- 
thophosphates. This  change  is  quickly  effected  by  heat- 
ing with  acids. 

Metaphosphoric  acid  is  formed  from  orthophosphoric 
acid  by  evaporating  to  complete  dryness  and  then  igniting 
strongly.  It  forms  a  solid  glassy  mass  and  is  commonly 
known  as  glacial  phosphoric  acid.  It  dissolves  readily  in 
water,  and,  as  explained,  this  solution  yields  the  ortho  acid 
on  boiling.  Solutions  of  metaphosphoric  acid  yield  white 
precipitates  with  egg  albumin  and  with  calcium  chloride. 
These  precipitates  do  not  form  with  the  ortho  and  pyro 
acids. 


GENERAL  CHEMISTRY.  191 

The  three  phosphoric  acids  may  be  looked  upon  as  re- 
lated to  each  other  in  this  manner  : 

P205  +  HaO  =  2HP08 

P205 

PO 


Orthophosphoric  acid  has  three  replaceable  hydrogen 
atoms,  or  is  tribasic;  it  forms  therefore  three  classes  of  salts. 

H3PO4        is  phosphoric  acid. 
H2NaPO4  is  dihydrogen  sodium  phosphate. 
HNa2PO4  is  hydrogen  disodium  phosphate. 
Na3PO4      is  trisodium  phosphate. 

The  salt  HNa2PO4-f-12H2O  is  known  as  common  so- 
dium phosphate.  It  is  used  in  the  laboratory  as  a  test 
reagent. 

Tests  for  Phosphoric  Acid.  Orthophosphoric  acid 
and  the  orthophosphates  are  recognized  by  several  tests. 
They  yield  a  yellow  precipitate  with  silver  nitrate,  a  white 
crystalline  precipitate  with  an  alkaline  solution  containing 
ammonia,  ammonium  chloride  and  magnesium  chloride, 
and  finally  a  very  characteristic  .yellow  precipitate  with  a 
solution  of  ammonium  molybdate.  This  is  a  delicate  test 
and  will  be  here  illustrated. 

Ex.  131.  Take  a  few  cubic  centimeters  of  a  phosphate  solution  in 
a  test-tube,  add  several  drops  of  strong  nitric  acid  and  then  some  solu- 
tion of  ammonium  molybdate.  Allow  the  test-tube  to  stand  at  rest  ten 
minutes.  A  fine  yellow  crystalline  precipitate  settles  out.  This  pre- 
cipitate is  soluble  in  ammonia  water,  as  can  be  seen  by  trial.  The  com- 
position of  this  crystalline  substance  is  complex,  but  is  probably 
(NH4)3(Mo03)12  P04. 

The  meta  and  pyrophosphates  in  nitric  acid  solution, 
after  heating,  yield  exactly  the  same  test. 

PHOSPHORUS  AND  THE  HALOGENS. 

Phosphorus  forms  important  compounds  with  chlorine, 
bromine  and  iodine  by  direct  union. 


192  GENERAL  CHEMISTRY. 

Phosphorus  Trichloride  is  a  colorless  liquid  obtained 
by  passing  dry  chlorine  gas  over  phosphorus  heated  in  a 
retort.  The  phosphorus  burns  in  the  chlorine  atmosphere 
and  the  product  distills  over  into  a  dry,  cool  receiver. 
The  trichloride  boils  at  76°.  It  fumes  in  the  air  and  is 
decomposed  by  water. 


The  reaction  just  illustrated  is  one  of  very  great  impor- 
tance in  the  decomposition  of  bodies  which,  like  water, 
contain  the  hydroxyl  group,  OH.  In  organic  chemistry  the 
chlorides  and  bromides  of  phosphorus  find  extended  appli- 
cations in  the  preparation  of  halogen  derivatives  of  acids 
and  alcohols,  which  contain  the  OH  group.  In  illustra- 
tion, the  formula  of  common  alcohol  is  C2H5OH,  and  this 
combines  with  PC13  or  PBr3  in  this  manner: 

3C2H5OH  +  PCl3=3C2H5Cl-fH303P, 
or,  expressed  graphically, 

C2H5OH         (  Cl     C2H5C1     HO  ) 
-   C1  = 


C2H5OH+P-{  C1  =  C2H5C1  +  HO  VP 
C2H5OH         (  Cl     C2H5C1     HO  j 

Acetic  acid  has  the  composition  C2H3OOH.  This 
reacts  with  the  trichloride  as  follows: 

3C2H3OOII  +  PCl3=:3C2H3OCl-hH3O3P, 

acetyl  chloride  and  phosphorous  acid  being  formed. 

The  student  should  note  these  typical  reactions  care- 
fully, as  he  will  have  occasion  to  use  them  in  the  study  of 
the  complex  compounds  of  carbon.  The  pentachloride  of 
phosphorus,  described  below,  is  often  used  in  preference 
to  the  trichloride  because  of  its  more  energetic  action. 

Phosphorus  Pentachloride  is  made  by  the  continued 
action  of  chlorine  on  the  trichloride.  The  chlorine  is 
absorbed  and  a  dry  solid  mass  results.  This  is  the  crude 
pentachloride  which  may  be  purified  by  sublimation.  Like 


GENERAL  CHEMISTRY.  193 

the  trichloride  it  is  decomposed  by  water,     forming  the 
oxychloride  first: 

PCl5+H20=:2HCl-fPOCl3. 
An  excess  of  water  reacts  in  this  manner: 


Phosphorus  Bromides.  Phosphorus  and  bromine 
combine  directly,  forming  the  tribromide,  PBr3,  which  is  a 
colorless  liquid,  boiling  at  about  175°,  and  having  a  spe- 
cific gravity  of  2.925.  Phosphorus  pentabromide  is  a  yel- 
low crystalline  solid,  formed  by  the  action  of  bromine  on 
the  tribromide. 

Phosphorus  Iodides.  Two  iodides  are  known,  ob- 
tained by  mixing  solutions  of  iodine  and  phosphorus  in 
carbon  disulphide.  After  the  combination  is  complete  the 
disulphide  may  be  distilled  off,  leaving  the  iodides  as  crys- 
talline products.  By  using  a  certain  amount  of  iodine  the 
iodide,  P2I4,  is  obtained,  while  with  a  larger  proportion 
the  tri-iodide,  PI3,  is  made.  The  latter  resembles  the  cor- 
responding chlorine  and  bromine  compounds  in  its  behav- 
ior with  water: 

PI3-f3H2O=:3HI+H3PO3. 

Advantage  is  taken  of  this  reaction  in  the  preparation 
of  hydriodic  acid. 

Two  compounds  with  fluorine,  PF3  and  PF5,  are 
known,  as  are  also  several  compounds  with  nitrogen,  oxy- 
gen and  hydrogen.  But  none  of  these  may  be  considered 
important. 

ARSENIC. 

Occurrence.  This  element  is  rather  widely  distrib- 
uted in  nature,  being  found  in  many  ores,  and  also  in  the 
free  state.  The  important  ores  which  contain  it  are 
realgar,  As2S2,  orpiment,  As2S3,  arsenical  pyrite, 
Fe2S2As,  arsenical  nickel,  NiAs,  and  arsenical  iron,  FeAs2. 


194  GENERAL  CHEMISTRY. 

Much   of    the    arsenic  of   commerce    is  obtained  as  a  by- 
product in  the  working  of  these  ores. 

History.  The  sulphides  of  arsenic  have  been  known 
from  the  earliest  times  and  the  alchemists  were  acquainted 
with  the  preparation  of  the  white  oxide  (commonly  called 
arsenic)  from  these  ores.  Albertus  Magnus,  in  the  twelfth 
century,  first  stated  that  a  metal-like  substance  is  con- 
tained in  this  white  arsenic,  but  this  view  was  not  com- 
monly admitted  until  near  the  end  of  last  century  when  it 
was  experimentally  found  that  the  white  body  is  the  oxide 
of  arsenic  or  arsenicum. 

Preparation.  What  is  known  as  metallic  arsenic  is 
usually  made  by  the  sublimation  of  arsenical  pyrite: 

Fe2S2As  =  2FeS-f  As. 

It  may  also  be  obtained  by  heating  the  white  oxide 
with  powdered  coal  or  charcoal. 

Properties.  When  prepared  by  reduction,  arsenic  is 
a  dark  steel  gray,  brittle  mass,  with  metallic  luster.  It  is 
easily  sublimed,  yielding  a  yellow  vapor,  and  at  a  high 
temperature  it  combines  readily  with  oxygen  to  form  the 
trioxide.  In  moist  air  it  oxidizes  at  the  ordinary  temper- 
ature to  form  the  same  oxide.  The  specific  gravity  of  ar- 
senic is  5.73,  and  in  this  and  other  properties  it  resembles 
the  true  metals.  In  its  combinations,  however,  it  acts  usu- 
ally as  a  nonmetallic  element.  It  alloys  with  some  of  the 
metals.  It  hardens  lead,  and  is  therefore  frequently  added 
to  the  latter  element  in  the  manufacture  of  shot. 

ARSENIC  AND  HYDROGEN. 

Arsenic  forms  two  compounds  with  hydrogen,  one 
of  which,  AsH3,  is  a  gas,  while  the  other,  As2H2,is  a  solid. 
The  first  only  is  important.  A  compound  having  the 
formula  AsH,  probably,  has  also  been  described. 

Hydrogen  Arsenide  or  Arsinei  This  gas  was  pre- 
pared by  Scheele  in  pure  condition  in  1775,  and  later  by 


GENERAL  CHEMISTRY.  195 

Proust.  It  is  always  formed  when  nascent  hydrogen  is 
liberated  in  presence  of  an  arsenic  compound  in  acid  solu- 
tion, and  as  usually  liberated  is  mixed  with  a  large  excess 
of  hydrogen.  It  may  be  made  in  pure  condition  by  the 
action  of  sulphuric  acid  on  zinc  arsenide: 

Zn3As2+3H2SO4=3ZnSO4+2AsH3. 

The  gas  is  intensely  poisonous  and  one  of  the  early  ex- 
perimenters, Gehlen,  lost  his  life  by  inhaling  a  very  small 
quantity. 

It  is  easily  made  in  impure  form  by  the  addition  of  a 
solution  of  any  arsenic  compound,  best  neutral  or  slightly 
acid,  to  a  hydrogen  generator  in  action.  The  nascent  hy- 
drogen decomposes  the  compound,  as  illustrated  by  this 
equation: 

As2O3+HH2=3H2O-f2AsH3. 

The  AsH3  escapes  mixed  with  the  excess  of  hydrogen 
given  off.  The  student  should  make  the  following  experi- 
ment, but  in  a  good  fume  closet  or  in  a  good  draught  of  air. 

In  no  case  should  more  than  the  minute  amount  of  ar- 
senic compound  referred  to  be  taken  for  experiment. 

Ex.  132.  Arrange  a  hydrogen  generator,  as  already  explained  and 
as  further  illustrated  in  the  figure  below.  To  the  generator  is  attached 
a  drying  tube  with  calcium  chloride,  and  following  this  is  a  tube  of  mod- 
erately hard  lead-free  glass,  narrowed  slightly  in  the  middle  and  drawn 
to  a  small  orifice  at  the  outer  end.  Charge  the  bottle  with  pure  zinc, 
free  from  arsenic,  and  add  pure  dilute  sulphuric  acid  in  the  usual  man- 
ner. Allow  the  evolution  of  gas  to  continue  until  the  air  is  all  expelled, 
and  then  light  it  at  the  drawn  out  end  of  the  delivery  tube.  The  burning 
gas  is  pure  hydrogen  only,  if  the  materials  used  are  pure.  In  any  event, 
test  the  gas  by  holding  a  clean  cold  porcelain  dish  against  the  flame  sev- 
eral minutes.  If  after  five  minutes  no  black  stain  or  deposit  appears  on 
the  dish  the  gas  may  be  considered  pure  for  the  present  purpose.  Now 
prepare  a  very  dilute  solution  of  sodium  arsenite,  or  arsenous  acid,  and 
pour  5  drops  of  this  through  the  funnel  tube  into  the  generator.  Hold 
the  porcelain  dish  again  in  the  flame  and  observe  that  after  a  time  a 
shining  black  deposit  of  arsenic  collects.  Allow  several  of  these  stains 
or  deposits  to  form  on  different  parts  of  the  dish.  (The  stains  may  be 
further  examined  as  will  be  explained  shortly.)  Next  heat  the  delivery 
tube  with  the  Bunsen  flame  a  short  distance  in  front  of  the  central  nar- 
rowed part,  that  is  between  the  generator  and  the  constriction.  At  a 
high  temperature  the  hydrogen  arsenide  decomposes,  yielding  arsenic, 


196 


GENERAL  CHEMISTRY. 


which  precipitates  on  the  cooler  part  of  the  tube,  and  hydrogen,  which 
burns  at  the  end.  If  the  heat  is  applied  at  the  right  point  the  arsenic 
deposit  forms  in  the  contracted  portion  of  the  tube  where  it  can  be  read- 
ily seen.  The  liberated  arsenic,  being  slightly  volatile,  is  deposited 
some  distance  beyond  the  point  most  strongly  heated.  The  stains  in 
the  dish  may  now  be  tested.  Pour  in  a  few  drops  of  a  fresh  solution  of 
sodium  hypochlorite  and  allow  thisto  flow  around  the  dish.  The  deposit 
goes  into  solution  almost  instantly.  This  behavior  distinguishes  arsenic 
from  very  similar  antimony  stains  to  be  described  later.  In  performing 
this  experiment  care  must  be  taken  to  have  the  corks  and  fittings  of  the 
apparattis  perfectly  light,  and  to  keep  the  gas  burning^  after  the  addi- 
tion of  the  arsenic  solution. 

The  experiment  just  described  constitutes  what  is  called 


FIG.  21. 


the  Marsh  test  for  arsenic.  It  is  exceedingly  delicate,  and 
further  details  concerning  it  may  be  found  in  books  on 
qualitative  analysis,  as  it  is  practically  important. 

Hydrogen  arsenide  is  not  soluble  in  water  to  a  great 
extent.  One  volume  of  cold  water  will  take  up  about  five 
volumes  of  the  gas,  but  the  product  decomposes  on  stand- 
ing. In  the  above  experiment  it  forms  water  and  arsenous 
oxide  on  burning. 

2AsH3-f3O2=As2O3-f3H2O. 
The  white  oxide  imparts  a  light  color  to  the  flame.     If 


GENERAL  CHEMISTRY.  197 

the  flame  is  much  cooled,  as  happens  when  the  cold  por- 
celain is  held  in  it,  the  hydrogen  alone  burns,  forming 
water,  while  the  arsenic  is  deposited  on  the  cold  surface  in 
the  free  state,  or  possibly  as  solid  hydrogen  arsenide, 
AsH. 

ARSENIC    AND    OXYGEN. 

Arsenic  forms  two  compounds  with  oxygen,  the  trioxide 
or  arsenous  oxide,  As2O3,  and  the  pentoxide,  As2O5. 
Both  are  well  known  and  common. 

Arsenous  Oxide.  This  is  the  body  commonly  known 
as  white  arsenic.  It  appears  in  commerce  in  several 
forms,  as  fine  powder,  granular  or  massive.  The  massive, 
or  vitreous,  variety  is  formed  by  the  sublimation  of  the 
crude  white  powder,  which  is  a  by-product  in  the  roasting 
of  certain  ores  containing  arsenides.  Arsenous  oxide  has 
the  composition  As2O3,  or  As4O6  probably.  It  is  slightly 
soluble  in  water,  the  solution  containing  the  hydroxide, 
As(OH)3,  or  arsenous  acid. 

Arsenous  oxide  is  a  violent  poison,  this  depending  on 
its  solubility  in  the  liquids  of  the  stomach  or  intestines.  It 
is  easily  reduced  to  the  free  or  metallic  condition,  as  illus- 
trated by  the  following  experiment : 

Ex.  133.  Cut  from  a  piece  of  glass  tubing  having  an  internal  di- 
ameter of  about  3  to  4  Mm.,  a  length  of  20  Cm.  Hold  this  in  the  flame 
until  it  is  softened  in  ths  middle,  then  draw  it  out  about  4  Cm.  and 
melt  until  the  two  pieces  separate,  thus  making  two  closed  tubes, 
pointed  and  sealed  at  one  end.  Drop  a  minute  amount  of  arsenous  ox- 
ide into  one  of  these  tubes  and  tap  until  the  powder  collects  in  the  point. 
Then  drop  in  a  small  fragment  of  charcoal,  which  is  just  large  enough 
to  fill  the  wide  part  of  the  tube,  but  which  cannot  follow  the  oxide  into 
the  extreme  point.  Now  heat  the  tube  in  the  flame,  the  part  on  which 
the  charcoal  rests  first,  and  then  the  pointed  end  below  it.  The  white 
oxide  is  volatilized  and  passing  up  comes  in  contact  with  the  hot  char- 
coal, suffering  reduction  to  the  free  condition  as  expressed  in  this 
equation: 

As2O3-f3C=As2-f3CO. 

The  liberated  arsenic  collects  as  a  dark  mirror  on  the  colder  part 
of  the  tube  above  the  charcoal.  Allow  the  tube  to  cool,  break  off  the 
lower  end  and  heat  it  again  carefully.  The  current  of  hot  air  passing 
over  the  arsenic  oxidizes  it  to  arsenous  oxide,  and  this  forms  a  white 
ring  higher  up  if  the  heat  is  not  too  strong. 


198  GENERAL  CHEMISTRY. 

On  the  large  scale  the  reduction  is  effected  in  iron 
retorts. 

Arsenous  oxide,  as  explained,  is  but  slightly  soluble  in 
water.  In  dilute  hydrochloric  acid  it  dissolves,  but  does 
not  form  the  chloride.  Heated  with  nitric  acid  oxidation  to 
arsenic  oxide  takes  place.  Alkali  solutions  dissolve  it 
readily,  forming  arsenites,  as  illustrated  by  this  equation: 

As8Os+6NaOH  =  2Na8AsOs+3H8O. 

The  subject  of  the  solubility  of  arsenous  oxide  in  water 
and  other  liquids  is  a  very  important  one  from  a  toxico- 
logical  standpoint,  and  much  has  been  written  on  it.  The 
solubility  of  the  pure  oxide  has  been  found  to  vary  much 
with  its  physical  condition. 

The  Arsenites.  Some  of  these  are  common  sub- 
stances. Those  of  the  alkali  metals  are  readily  soluble  in 
water  and  are  very  poisonous.  The  arsenites  of  silver, 
copper  and  some  other  metals  are  insoluble  in  water  and 
appear  therefore  when  a  soluble  arsenite  solution  is  added 
to  a  solution  of  copper  sulphate,  silver  nitrate,  etc.  The 
silver  arsenite  is  yellow;  copper  arsenite  is  green.  The 
best  known  of  the  arsenites  is  sodium  arsenite,  Na3AsO3. 
It  is  employed  for  many  purposes,  especially  in  the  produc- 
tion of  the  insoluble  arsenites  and  of  embalming  fluids. 
Some  of  the  insoluble  arsenites  have  been  largely  used  as 
pigments,  but  now,  fortunately,  they  are  being  displaced 
by  the  bright  and  harmless  coal-tar  dyes.  Paris  green  is 
a  crude  product  made  by  precipitating  sodium  or  potas- 
sium arsenite  by  copper  acetate.  It  is  essentially  copper 
aceto-arsenite.  but  of  somewhat  variable  composition.  It 
is  known  also  as  Schweinfurth  green,  emerald  green  and 
French  green.  Scheele's  green  is  copper  arsenite,  CuHAsO3. 
All  of  these  compounds  are  poisons  and  should  not  be 
used. 

Arsenic  Oxide,  As2O5.  This  product  is  easily  made 
by  heating  the  trioxide  with  strong  nitric  acid,  which  acts 
as  an  oxidizing  agent. 


GENERAL    CHEMISTRY.  199 

As  the  oxide  is  soluble,  it  is  left  after  the  action  in  the 
form  of  arsenic  acid,  H3AsO4.  On  evaporation  and  heat- 
ing to  a  temperature  of  about  300°  this  acid  decomposes 
and  leaves  the  pure  anhydrous  pentoxide.  At  a  red  heat 
it  breaks  up  into  As2O3-[-O2.  The  pentoxide  is  much  more 
soluble  in  water  than  the  trioxide,  and  forms  a  strong  solu- 
tion of  arsenic  acid. 

Arsenic  Acid.  This  acid,  as  just  explained,  has  the 
composition  H3AsO4,  and  is  best  made  by  boiling  the  tri- 
oxide with  nitric  acid.  It  appears  in  commerce  as  a  thick 
liquid  with  a  strong  acid  taste,  from  which  a  crystalline 
hydrated  solid  may  be  obtained  by  cooling  to  a  low 
temperature.  At  100°  the  water  of  crystallization  is  lost 
and  the  true  acid,  H3AsO4,  remains,  while  at  a  higher 
temperature  pyroarsenic  acid,  H4As2O7,  and  finally  met- 
arsenic  acid,  HAsO3,  result.  These  bodies  correspond  to 
pyro  and  metaphosphoric  acids. 

Arsenic  acid  parts  with  oxygen  readily  enough  to  serve 
as  a  good  oxidizing  agent  and  it  is  so  employed  in  the  man- 
ufacture of  aniline  red  on  the  large  scale. 

The  Arsenates.  In  its  behavior  with  bases  arsenic 
acid  closely  resembles  phosphoric  acid  and  forms  three 
classes  of  salts,  as  Na3AsO4,  Na2HAsO4  and  NaH2AsO4. 
The  salts  of  arsenic  acid  are  isomorphous  with  those  of  phos- 
phoric acid.  With  silver  nitrate  the  arsenates  yield  a  red 
precipitate,  Ag3AsO4,  while  the  arsenites  yield  a  yellow 
precipitate,  Ag3AsO3. 

ARSENIC  AND  THE  HALOGENS. 

Compounds  of  arsenic  with  chlorine,  bromine  and 
iodine  are  known.  The  chloride  is  a  heavy  liquid  best 
made  bypassing  chlorine  gas  over  heated  arsenic;  the 
bromide  is  a  solid,  made  by  adding  powdered  arsenic  to  a 
solution  of  bromine  in  carbon  disulphide;  the  iodide  is  also 
a  solid  and  is  obtained  from  iodine  and  arsenic  in  the  same 
manner.  These  three  compounds  decompose  with  water : 

2AsCl3-f3H2O^As2O3 


200  GENERAL  CHEMISTRY. 

To  decompose  the  chloride  in  presence  of  hydrochloric 
acid  a  considerable  excess  of  water  is  necessary. 

ARSENIC  AND  SULPHUR. 

Two  sulphides  of  arsenic  occur  in  nature,  realgar, 
As2S2,  and  orpiment,  As2S3.  They  may  be  both  made  ar- 
tificially. Artificial  realgar  comes  into  commerce  as  a  ruby 
red  glassy  mass,  sold  as  a  pigment  and  as  a  depilatory  for 
use  of  tanners.  It  is  made  by  subliming  mixtures  of  ar- 
senical and  common  pyrites.  Commercial  orpiment  is 
known  as  King's  yellow,  and  is  made  by  sublimation  of  a 
mixture  of  sulphur  and  arsenous  oxide.  A  much  purer 
product  is  made  by  precipitating  a  solution  of  an  arsenite 
with  hydrogen  sulphide.  A  pentasulphide,  As2S5,  is  also 
known  as  a  yellow  solid  made  by  fusing  the  trisulphide 
with  sulphur.  Classes  of  salts  known  as  sulpharsenites 
and  sulpharsenates  exist,  corresponding  to  arsenites  and 
arsenates.  K3AsS3  and  K3AsS4  are  illustrations. 

Antidotes  in  Arsenical  Poisoning.  The  several  anti- 
dotes recommended  to  be  given  in  cases  of  arsenical  poison- 
ing are  substances  which  form  insoluble  compounds  with 
arsenic  acids  and  salts.  Probably  the  best  of  these  is 
freshly  precipitated  ferric  hydroxide,  which  must  be  well 
washed  with  water  and  administered  in  relatively  large 
quantity.  With  arsenous  acid  or  arsenites  it  forms  an 
insoluble  basic  arsenite  of  iron.  Magnesia  is  sometimes 
used  for  the  same  purpose,  and  a  product  made  by  agita- 
tion of  an  excess  of  magnesia  with  ferric  sulphate  solution 
has  also  given  good  results.  This  contains  ferric  hydroxide 
and  magnesium  sulphate  along  with  magnesium  oxide  and 
hydroxide.  These  antidotes  were  proposed  primarily  for 
arsenous  oxide,  but  they  act  equally  well  with  most  of  the 
other  compounds. 

Tests  for  Arsenic. 

Because  of  the  great  importance  of  the  subject  much 
attention  has  been  given  to  the  tests  for  arsenic.  One  of 


GENERAL  CHEMISTRY.  201 

the  best  of  these  for  the  recognition  of  traces  is  the  Marsh 
test  described  above.  Another  is  the  Reinsch  test,  which 
may  be  illustrated  as  follows  : 

Ex.  134.  Pour  10  cubic  centimeters  of  pure  diluted  hydrochloric 
acid  into  a  test-tube  and  add  a  bit  of  bright  pure  copper  foil.  A  piece 
a  centimeter  square  will  answer.  Boil  the  acid  a  minute  or  two  and 
observe  that  the  foil  remains  bright,  which  must  be  the  case  in  absence 
of  arsenic.  Now  add  a  few  drops  of  a  dilute  arsenical  solution  and  heat 
again  to  boiling.  A  greyish  black  deposit  forms  on  the  copper  and  this 
consists  of  an  alloy  of  copper  and  arsenic.  This  deposit  may  be 
obtained  from  extremely  dilute  solutions;  the  test  is  therefore  a  very 
delicate  one. 

Arsenites  and  arsenates  are  often  recognized  by  their 
behavior  with  silver  solution,  and  all  arsenic  compounds 
yield  a  precipitate  with  hydrogen  sulphide  in  properly  pre- 
pared solutions.  The  precautions  to  be  taken  in  applying 
these  tests  are  explained  in  books  on  qualitative  analysis. 


CHAPTER  X. 


CARBON    AND    SOME    OF     ITS     IMPORTANT     COM- 
POUNDS. 

WE  HAVE  here  a  very  abundant,  and  at  the  same  time 
one  of  our  most  important  elements,  which  is  found 
very  widely  distributed  in  nature,  in  vegetable,  animal  and 
mineral  substances,  and  also  in  the  free  state.  Carbon 
occurs  in  the  three  distinct  forms  of  diamond,  graphite  and 
amorphous  carbon.  Of  the  last  there  are  several  varieties, 
as  charcoal,  coke,  lampblack  and  boneblack. 

History.  Charcoal,  graphite  and  the  diamond  have 
been  known  from  remote  antiquity,  but  that  they  are  all 
varieties  of  carbon  was  not  recognized  until  within  com- 
paratively recent  times.  In  the  seventeenth  century  it  was 
shown  that  the  diamond  may  be  caused  to  disappear  when 
placed  in  the  focus  of  a  large  double  convex  lens  in  bright 
sunlight.  That  this  disappearance  is  due  to  combustion 
at  the  expense  of  the  oxygen  of  the  air  was  shown  a  hun- 
dred years  later  by  Lavoisier  and  others,  who  demonstrated 
that  carbon  dioxide  is  produced  in  the  combustion. 
Scheele  showed  somewhat  earlier  that  graphite  may  be 
converted  by  oxidation  into  carbon  dioxide,  and  between 
1796  and  1800  it  was  demonstrated  that  equal  weights  of 
charcoal,  graphite  and  diamond  yield  the  same  amount  of 
this  gas. 

THE  DIAMOND. 

This  variety  of  carbon  is  found  crystallized  in  several 
forms  belonging  to  the  regular  system.  Octahedra  are  the 
most  common.  For  hundreds  of  years  attempts  have  been 


GENERAL  CHEMISTRY.  203 

made  to  produce  the  diamond  artificially,  but  always 
without  success  until  quite  recently,  when  Moissan 
succeeded  in  producing  very  small  crystals  by  sub- 
jecting molten  cast  iron  containing  carbon  to  very 
great  pressure.  This  pressure  was  secured  by  the  rapid 
chilling  of  the  surface  of  the  iron,  from  which  contraction 
of  the  chilled  layer  on  the  still  liquid  center  followed.  The 
crystals  so  produced  are  very  small  and  have  no  commer- 
cial value,  but  the  process  is  one  of  great  scientific  inter- 
est and  may  lead  to  something  more. 

Regular  and  clear  forms  of  the  natural  diamond  are 
highly  prized  as  gems,  while  dark  and  opaque  crystals  are 
used,  because  of  their  great  hardness,  in  pointing  rock 
drills,  and  in  powder  for  grinding  the  gems.  The  dia- 
mond is  the  hardest  substance  known.  Its  specific  grav- 
ity is  3.5  to  3.6.  When  heated  to  a  sufficiently  high  tem- 
perature in  oxygen  the  diamond  burns  to  carbon  dioxide, 
leaving  only  a  trace  of  ash. 

GRAPHITE. 

This  form  of  carbon  is  known  also  as  plumbago  or 
black  lead.  It  is  found  in  many  localities,  but  the  largest 
deposits  are  in  Ceylon  and  Siberia.  There  are  also  very 
considerable  deposits  in  the  United  States. 

Properties.  Although  a  variety  of  carbon,  this  sub- 
stance oxidizes  with  extreme  difficulty.  When  heated  in 
the  air  or  oxygen  it  undergoes  but  little  change,  but  certain 
oxidizing  agents  attack  it  at  a  high  temperature.  It  has  a 
specific  gravity  of  2.25  in  pure  form  and  is  a  moderately 
good  conductor  of  heat  and  electricity. 

Uses.  Because  of  its  very  imperfect  combustibility 
and  infusibility  graphite  is  largely  used  in  the  manufacture 
of  the  so-called  plumbago  or  black  lead  crucibles  employed 
in  many  industries,  especially  in  the  production  of  crucible 
steel.  The  graphite  for  this  purpose  is  mixed  with  suffi- 
cient clay  to  make  it  coherent.  It  is  used  in  our  common 
"lead"  pencils  and  as  a  protective  coating  for  iron;  com- 


204  GENERAL  CHEMISTRY. 

mon  stove  polish  is  usually  fine  graphite.  It  is  frequently 
employed  as  a  coating  for  gunpowder  to  render  the  grains 
impervious  to  moisture,  and  because  of  its  un'ctuous  prop- 
erty is  often  used  as  a  lubricant  instead  of  oil. 

AMORPHOUS  CARBON. 

Charcoal,  coke,  lampblack,  boneblack  and  hard  coal  are 
the  important  varieties  of  amorphous  carbon.  Charcoal 
and  coke  are  distinguished  by  their  combustibility,  lamp- 
black by  its  properties  as  a  pigment,  and  boneblack  as  a 
filtering  material  for  the  clarification  of  many  organic 
liquids.  The  preparation  of  charcoal  and  coke  will  be 
illustrated  by  experiments. 

Ex.  135.  Charge  a  small  retort,  consisting  of  a  piece  of  iron  gas 
pipe  20  Cm.  long  and  15  Mm.  in  internal  diameter,  and  closed  air-tight 
at  one  end  by  an  iron  plug,  with  shavings  or  fine  splinters,  and  attach  it 
to  a  delivery  tube  and  receiver  as  shown  in  the  figure.  Heat  strongly 
with  a  Bunsen  burner.  The  wood  is  decomposed,  and  a  mixture  of 
water,  acid  liquids  and  gas  passes  over.  The  gas  collects  in  the  bottle, 
while  the  other  substances  condense.  When  the  evolution  of  the  gas 
has  ceased,  withdraw  the  delivery  tube  and  then  remove  the  lamp. 
After  the  iron  tube  has  cooled  shake  out  the  contents,  or  remove  them 
with  a  bit  of  wire,  and  test  by  burning.  Close  the  mouth  of  the  bottle 
with  a  glass  plate,  invert  it  and  open  to  the  air,  and  apply  a  match.  The 
gas  burns  with  a  bright  flame. 

The  process  illustrated  here  is  known  as  destructive 
distillation.  Wood  is  a  highly  complex  substance,  contain- 
ing a  large  amount  of  carbon.  If  heated  in  the  air  it  burns 
— that  is,  it  combines  with  the  oxygen  and  is  wholly 
consumed,  yielding  only  gases  and  water  vapor,  with  a 
minute  trace  of  ash.  But  when  heated  to  a  high  tem- 
perature away  from  the  air,  as  in  the  above  experiment,  a 
very  complicated  reaction  takes  place,  in  which  many  prod- 
ucts are  formed.  Besides  carbon,  the  wood  contains  hy- 
drogen and  oxygen  in  large  amount,  and  a  little  nitrogen. 
These  elements  combine  with  each  other  and  with  a  part 
of  the  carbon  to  form  new  compounds,  but  the  rest  of  the 
carbon  is  left  in  solid,  nonvolatile  form,  and  is  called  char- 
coal. A  part  of  the  liquid  which  distills  in  the  above  ex- 
periment forms  a  tarry  substance  on  cooling,  but  as  made 


GENERAL  CHEMISTRY. 


205 


here  the  amount  is  too  small  to  be  separated.  Acetic  acid, 
methyl  alcohol  or  wood  spirit,  and  acetone  are  other  very 
important  substances  present  in  the  products  of  the  distil- 
lation, and  on  the  large  scale  they  are  made  by  this  proc- 
ess in  quantity. 

The  gas  collected  in  the  bottles  is  called  wood  gas,  and 
at  one  time  was  made  for  illuminating  purposes  by  distill- 
ing wood  in  large  retorts. 

On  the  large  scale  in  the  manufacture  of  charcoal  wood 
is  heated  in  large  iron  or  brick  retorts,  or  it  may  be  cut  up 


FIG. 


into  cordwood  lengths,  stacked  in  circular  heaps,  covered 
with  a  thin  layer  of  earth  or  sods  and  ignited  at  openings 
left  at  the  base  of  the  heap.  A  hole  left  at  the  top  of  the 
heap  permits  a  slow  circulation  of  air  from  which  a  partial 
combustion  follows;  a  little  of  the  wood  is  consumed,  fur- 
nishing heat  enough  to  char  the  rest. 

Ex.  136.  Repeat  the  last  experiment,  using  soft  coal  in  small 
lumps  instead  of  wood.  Coke  remains  in  the  tube,  while  a  gas  similar  to 
ordinary  illuminating  gas  collects  in  the  bottle  and  can  be  tested  as 
before. 

In  the  dry  distillation  of  the  soft  coal,  the  gas  obtained 
burns  with  a  better  illuminating  flame  than  that  obtained 


206  GENERAL  CHEMISTRY. 

in  the  other  case.  Coal  differs  from  wood  in  containing 
more  carbon  and  much  less  oxygen,  from  which  it  follows 
that  the  products  of  the  distillation  of  coal  must  be  char- 
acterized by  a  lower  percentage  of  oxygen.  Some  of  these 
will  be  described  later;  we  are  at  present  concerned  with 
the  residues  left  in  the  retort.  We  have  here  two  valuable 
varieties  of  amorphous  carbon,  which  are  produced  on  the 
large  scale  in  enormous  quantities,  to  be  used  for  sev- 
eral purposes. 

Ordinary  charcoal,  when  freshly  heated,  has  a  marked 
power  of  absorbing  gases,  and  hence  is  sometimes  used  for 
purifying  contaminated  atmospheres.  It  absorbs  many 
solid  substances  also,  and  for  this  reason  is  employed  in 
the  construction  of  filters  for  the  filtration  of  water.  The 
value  of  charcoal  for  this  purpose  is  by  no  means  as  great 
as  was  at  one  time  supposed.  Charcoal  and  coke  are 
mainly  used,  however,  in  the  smelting  or  reduction  of  ores, 
that  is,  in  the  operations  in  which  certain  metals  are  freed 
from  the  substances  combined  with  them  in  their  ores. 
The  very  marked  affinity  of  carbon  for  oxygen  comes  in 
play  here. 

Boneblack,  or  bone  charcoal,  is  an  impure  form  of  car- 
bon left  in  the  dry  distillation  of  bones.  It  is  employed 
mainly  in  the  clearing  of  turbid  or  colored  organic  liquids, 
which  can  be  illustrated  by  this  experiment: 

Ex.  137.  Pour  some  water  strongly  colored  with  indigo  solution 
into  a  porcelain  dish  and  add  half  its  volume  of  boneblack.  Boil  the 
mixture  and  then  throw  it  on  a  paper  filter.  The  liquid  which  runs 
through  will  be  found  colorless  or  nearly  so.  The  experiment  may  be 
made  also  in  this  manner:  Place  a  paper  filter  in  a  small  funnel  and 
half  fill  it  with  the  boneblack.  Now  pour  the  liquid  through  and  repeat 
this  operation  several  times  until  the  color  is  removed. 

The  rapidity  of  the  absorption  of  coloring  substances  is 
greatly  accelerated  by  heat,  but  on  the  large  scale  colored 
sugar  syrups  and  other  liquids  are  decolorized  by  forcing 
them  through  long  columns  of  the  boneblack  packed  in 
iron  cylinders.  The  animal  charcoal  removes  not  only  the 
coloring  matter  but  many  other  objectionable  substances 
as  well.  This  form  of  carbon  is  frequently  used  in  the 


GENERAL  CHEMISTRY. 


207 


laboratory  for  the  purification  of  organic  substances  under 
preparation. 

Lampblack  is  the  carbon  which  separates  from  the 
smoke  produced  in  the  imperfect  combustion  of  gases  and 
vapors.  To  secure  the  lampblack  in  quantity  it  is  simply 
necessary  to  limit  the  supply  of  air  so  that  the  combustion 
must  take  place  slowly  and  incompletely.  Crude  turpen- 
tine, pine  knots,  rosin  and  natural  gas  are  all  utilized  in 
this  industry.  The  carbon,  being  deposited  from  the  con- 
dition of  a  gas  or  vapor,  and  not  left  as  a  residue  as  in  the 
case  of  charcoal,  boneblack  and  coke,  remains  in  an  ex- 
tremely fine  state  of  subdivision,  and  on  this  its  value  as  a 
pigment  largely  depends.  It  is  likewise  quite  insoluble 
in  water  or  other  liquid. 

Coal  exists  in  several  varieties  distinguished  mainly  by 
their  proportion  of  carbon.  All  kinds  of  coal  were  formed 
by  the  slow  decay  of  vegetation  in  absence  of  air,  and  in  this 
decomposition  the  original  vegetable  matter  lost  a  great  deal 
of  its  nitrogen,  hydrogen  and  oxygen,  leaving  a  more  or 
less  pure  carbon  residue.  Anthracite,  or  hard  coal,  is  the 
variety  richest  in  carbon;  bituminous  or  soft  coal  contains 
carbon  mixed  with  very  complex  hydrocarbons.  Many 
grades  of  soft  coal  are  known,  some  being  much  more  val- 
uable than  others.  Cannel  coal  may  be  mentioned  here  as 
a  soft  coal  rich  in  volatile  matter  which  burns  with  a 
bright,  yellow  flame.  The  following  short  table  gives  the 
average  composition  of  the  important  kinds  of  coal.  The 
original  wood  is  added  for  comparison,  and  peat  also,  be- 
cause it  is  a  stage  in  the  formation  of  many  coals.  The 
ash  is  deducted  in  the  analyses. 


CARBON. 

HYDROGEN. 

OXYGEN. 

NITROGEN. 

Wood 

50  2 

6    1 

43.7 

Peat          

61  5 

5.7 

32.8 

Lignite. 

67  9 

5.7 

25.4 

1.0 

Bituminous  .... 
Cannel  

79.4 
83  0 

5.3 
7.5 

13.9 

8.5 

1.4 
1.0 

Anthracite  

92.4 

3.5 

3.0 

1.1 

208  GENERAL  CHEMISTRY. 

When  anthracite  is  burned  little  beyond  carbon  diox- 
ide is  formed,  while  with  wood  and  the  softer  kinds  of 
coal  a  great  deal  of  water  is  produced  also. 

CARBON  AND  OXYGEN. 

The  combination  of  carbon  with  oxygen  has  been  illus- 
trated in  our  experiments  on  oxygen.  Two  different  com- 
pounds of  these  elements,  both  gases,  are  known,  one 
called  carbon  monoxide,  and  represented  by  the  formula 
CO,  while  the  other  is  called  carbon  dioxide  and  is  repre- 
sented by  the  formula  CO2.  The  second  is  the  most  im- 
portant and  will  be  described  first. 
• 

Carbon  Dioxide. 

While  this  substance  is  easily  made  by  the  combustion 
of  carbon  in  oxygen  it  is  best  prepared  for  laboratory  uses 
by  the  decomposition  of  a  carbonate,  or  combination  of  car- 
bon and  oxygen  with  a  metal.  The  most  convenient  car- 
bonate we  can  use  for  this  purpose  is  limestone,  or  the 
allied  substance,  marble. 

Ex.  138.  Arrange  the  apparatus  as  used  in  the  production  of  hy- 
drogen sulphide  and  charge  it  with  small  pieces  of  soft  marble  or  lime- 
stone. Pour  water  through  the  funnel  tube  and  then  some  hydrochloric 
acid.  The  delivery  tube  leads  into  a  dry,  empty  bottle.  As  the  product 
is  a  gas  much  heavier  than  the  air,  it  can  be  readily  collected  by  dis- 
placement in  this  manner.  The  acid  acts  on  the  marble  or  limestone 
immediately,  decomposing  it  and  leaving  a  compound,  called  calcium 
chloride,  in  solution.  The  gaseous  carbon  dioxide  escapes  by  the  tube 
and  enters  the  bottle,  forcing  out  the  air.  Collect  several  bottles  of  the 
gas,  and  then  put  under  the  delivery  tube  a  bottle  containing  about  50 
Cc.  of  distilled  water,  to  which  about  an  equal  volume  of  clear  lime- 
water  is  added.  Allow  the  gas  to  bubble  into  this  while  that  in  the 
bottle  is  tested. 

Ex.  139.  Add  to  one  of  the  bottles  some  lime-water  and  shake 
thoroughly.  A  white  precipitate  forms  which  does  not  disappear  on 
shaking.  Dip  a  small  lighted  taper  into  a  second  bottle  and  notice  that 
it  is  speedily  extinguished.  By  means  of  a  piece  of  wire  hold  a  burning 
match  or  a  small  taper  in  the  bottom  of  a  clean  bottle  and  flour  into  this 
the  gas  from  a  third  collecting  bottle.  The  light  will  be  extinguished 
here  as  before.  In  a  fourth  bottle  plunge  some  burning  sulphur  in  a 
deflagrating  spoon. 


GENERAL  CHEMISTRY.  209 

By  examination  of  the  solution  left  in  the  generating 
bottle  and  the  amount  of  gas  obtained,  we  are  able  to  write 
this  equation  as  illustrating  the  decomposition  of  the 
marble: 

CaCO3+2HCl  =  CaCl24-CO2-fH2O. 

Calcium  Calcium 

carbonate  chloride 

One  hundred  parts  of  marble  yield  forty-four  parts  by 
weight  of  the  gas.  On  bringing  the  gas  into  contact  with 
the  lime-water  a  white  precipitate  forms  at  first,  as  shown 
by  one  of  the  experiments,  but  if  a  large  excess  of  the  gas  is 
passed  into  this  turbid  liquid  the  precipitate  formed  disap- 
pears almost  completely.  We  have  here  two  distinct  re- 
actions. In  the  first  a  certain  amount  of  the  gas  combines 
with  lime-water,  to  produce  a  precipitate  which  on  exam- 
ination is  found  to  be  identical  in  composition,  though  not 
in  external  appearance,  with  the  marble  used  in  making 
the  gas.  The  lime-water  contains  a  substance  known  as 
calcium  hydroxide,  about  which  more  will  be  said  later. 
The  combination  between  the  calcium  hydroxide,  CaO2H2, 
and  the  carbon  dioxide  is  represented  by  this  equation: 

CaO2H2+CO2=CaCO3-fH2O. 

Here,  as  in  other  reactions  we  have  had,  a  base  and  an 
acid  substance  unite  to  form  a  salt.  Calcium  carbonate  is 
decomposed  much  in  the  same  manner  in  which  potassium 
nitrate  and  sodium  chloride  were  decomposed  by  sulphuric 
acid. 

The  experiment  of  passing  an  excess  of  gas  into  the 
lime-water  mixture,  followed  into  its  second  stage,  leaves  a 
nearly  clear  solution.  A  simple  explanation  of  this  phe- 
nomenon may  be  given.  The  carbon  dioxide  in  presence 
of  water  acts  as  a  weak  acid,  which  finally  dissolves  the 
precipitated  calcium  carbonate,  making  a  new  soluble  sub- 
stance known  as  calcium  bicarbonate.  The  solution  in  the 
bottle,  therefore,  contains  this  substance.  Pour  out  some 
of  it  into  test-tubes  and  make  the  following  experiments: 

Ex.  140.  To  a  small  portion  of  the  solution  of  calcium  bicarbon- 
ate, just  described,  add  some  hydrochloric  acid.  Gas  bubbles  escape, 
showing  the  decomposition  of  the  substance  in  solution.  Boil  another 


210  GENERAL  CHEMISTRY. 

portion  in  a  test-tube  and  observe  that  gas  escapes  here  also,  while  a 
precipitate  forms.  In  this  case  heat  decomposes  the  bicarbonate,  driv- 
ing off  the  excess  of  absorbed  gas  and  leaving  the  ordinary  insoluble 
carbonate. 

These  experiments  illustrate  some  very  common  phe- 
nomena. Lake,  river  and  well  waters  contain,  often,  cal- 
cium bicarbonate,  formed  by  the  solvent  action  of  the  car- 
bon dioxide  of  the  air,  carried  down  by  the  rain  water,  on 
limestone  rocks.  Such  waters  are  called  hard.  When 
they  are  boiled  the  bicarbonate  is  decomposed  and  a  pre- 
cipitate of  carbonate  produced.  In  boilers  and  teakettles 
this  precipitate  collects  as  a  crust  or  scale. 

Carbon  dioxide  is  somewhat  soluble  in  water,  forming 
a  weak  solution  of  carbonic  acid.  It  is  absorbed  by  alkali 
solutions,  forming  carbonates. 

Ex.  141.  Lead  the  gas  from  the  generator  into  a  flask  containing 
about  150  Cc.  of  distilled  water.  After  fifteen  or  twenty  minutes  re- 
move this  flask  and  dip  the  delivery  tube  into  a  bottle  or  flask  contain- 
ing a  weak  solution  of  sodium  hydroxide.  Observe  that  the  gas  is  much 
more  perfectly  absorbed  here  than  in  the  water.  Of  the  aqueous  solu- 
tion now  pour  about  10  Cc.  into  a  test-tube  and  add  some  lime-water.  A 
precipitate  forms.  Place  the  flask  with  the  remainder  of  the  solution  on 
a  sand-bath  or  gauze,  heat  slowly  and  then  boil  ten  minutes.  During  the 
slow  heating  it  will  be  noticed  that  gas  bubbles  collect  along  the  sides  of 
the  flask  and  then  escape.  When  the  water  actually  boils  the  gas  bubbles 
pass  up  rapidly.  When  no  more  are  seen  to  escape  allow  the  water  to 
cool  and  to  some  of  it  add  lime-water  as  before.  A  precipitate  fails  to 
form  now,  showing  the  absence  of  carbon  dioxide.  Next  treat  a  portion 
of  the  solution  of  sodium  hydroxide,  into  which  the  gas  was  passed,  with 
hydrochloric  acid  ;  an  evolution  of  gas  is  produced.  Boil  the  remainder 
and  add  hydrochloric  acid  to  it.  We  notice  here,  also,  an  evolution  of 
gas,  showing  that  heat  does  not  drive  it  from  the  alkali  solution,  as  it 
did  from  the  water. 

The  tests  just  made  should  be  thoroughly  understood. 
We  illustrate  by  this  equation  the  formation  of  the  weak 
acid  produced  when  the  gas  is  passed  into  distilled 
water: 

C02+H20  =  H2C03. 

The  carbonic  acid,  like  sulphurous  acid  and  ammonium 
hydroxide,  is  not  stable,  but  is  decomposed  by  heat. 
Hence,  when  we  boil  its  aqueous  solution  it  soon  disap- 
pears. With  the  sodium  hydroxide,  on  the  other  hand,  it 


GENERAL  CHEMISTRY.  211 

forms  a  stable  compound  called  sodium  carbonate,  having 
the  composition  Na2CO3,  as  illustrated  in  this  manner  : 


This  sodium  carbonate  is  decomposed  by  hydrochloric 
acid  just  as  the  marble  dust  was  above.  Carbonates  in 
general  are  decomposed  by  hydrochloric  acid. 

Still  another  source  of  carbon  dioxide  must  be  exam- 
ined. It  has  been  shown  that  it  is  produced  in  combus- 
tion processes  where  carbonaceous  matter  is  burned.  Or- 
dinary wood  and  coal  burned  in  a  stove  or  furnace  give 
rise  to  this  gas.  Now  in  the  animal  body  a  kind  of  com- 
bustion is  in  progress  all  the  time,  and  as  a  result  of  this 
combustion,  as  of  all  others,  heat  is  given  off.  The  warmth 
of  the  body  is  derived  from  the  oxidation  or  burning  of  the 
substances  taken  as  food,  all  of  which  contain  carbon.  The 
most  important  product  of  this  combustion  of  the  food 
stuffs,  or  of  the  tissues  built  up  from  them,  is  carbon  diox- 
ide, and  this  leaves  the  body  by  the  agency  of  the  lungs  in 
the  breath,  or  expired  air.  We  may  test  for  it  and 
readily  recognize  it. 

Ex.  142.  Pour  some  clear  lime-water  in  a  beaker  and  blow  into  it 
through  a  glass  tube  dipping  beneath  its  surface.  The  liquid  becomes 
turbid  as  before.  After  the  action  has  been  continued  some  minutes 
add  hydrochloric  acid  to  the  contents  of  the  beaker.  The  liquid  be- 
comes clear,  while  gas  escapes. 

This  carbon  dioxide  thrown  off  from  the  lungs  plays  a 
very  important  part  in  nature.  From  this  source,  and  from 
the  combustion  of  fuel,  enough  enters  the  atmosphere  to 
make  about  3  volumes  in  every  10,000.  Were  it  not 
for  the  action  of  plants  in  absorbing  the  gas,  in  taking  up 
and  using  as  a  food  what  animals  reject,  the  air  would 
soon  become  unfit  for  breathing.  The  presence  of  the  gas 
in  the  air  can  be  readily  shown  by  exposing  a  little  clear 
lime-water  in  a  beaker  to  its  action  during  ten  or  fifteen 
minutes.  In  the  atmosphere  of  a  close,  occupied  room 
the  lime-water  becomes  turbid  very  soon.  In  the  outside 
atmosphere  a  longer  time  is  required. 

Carbon  dioxide  is  formed  also  in  the  process  known  as 


212  GENERAL  CHEMISTRY. 

alcohol  fermentation  which  will  be  fully  described  later. 
The  amount  of  the  gas  liberated  in  this  manner  is  enor- 
mous and,  until  recently,  was  allowed  to  go  to  waste.  In 
many  large  breweries  at  the  present  time  it  is  saved. 

Physical  Properties.  One  liter  of  carbon  dioxide  at 
standard  temperature  and  pressure  weighs  1.98  Gm.  Re- 
ferred to  air  it  has  a  specific  gravity  of  1.529.  At  tempera- 
tures below  30.9°  it  may  be  condensed  to  a  liquid  which 
has  a  specific  gravity  of  about  0.995  at  10°.  The  gas  is 
soluble  in  water,  one  volume  of  which  at  0°  dissolves  1.797 
volumes  of  the  gas,  while  at  15°  1.002  volumes  are  dis- 
solved and  at  20°  0.901  volumes. 

Uses.  Liquid  carbon  dioxide  is  now  a  common  arti- 
cle of  commerce.  It  is  sold  in  strong  steel  cylinders  and 
is  employed  for  charging  "soda."  fountains,  for  extinguish- 
ing fires  and  for  several  minor  purposes. 

The  Carbonates.  Carbon  dioxide  in  water  forms  car- 
bonic acid,  which  however  is  very  weak  and  unstable.  The 
salts  corresponding  are  the  common  carbonates,  some  of 
which  are  readily  decomposed  by  heat  with  loss  of  CO2, 
while  others  are  stable. 

Tests  for  Carbonates. 

All  carbonates  are  decomposed  by  mineral  acids  with 
evolution  of  CO2.  This  is  best  recognized  by  the  lime- 
water  test  given  above.  It  has  no  odor  and  in  this  respect 
also  differs  from  SO2.  Like  the  latter  gas,  it  possesses  the 
property  of  extinguishing  flame. 

Carbon  Monoxide. 

This  oxide  of  carbon  is  a  very  important  substance 
which  differs  from  the  other  in  essential  properties.  It  is 
not  a  natural  substance,  but  can  be  produced  by  several 
reactions,  one  of  which  is  here  given: 

Ex.  143.  Powder  10  or  15  Gm.  of  potassium  ferrocyanide, 
K4Fe(CN)6,  place  it  in  a  flask  of  300  Cc.  capacity,  having  a  funnel  tube 


GENERAL  CHEMISTRY.  213 

and  delivery  tube.  Add  ten  times  its  weight  of  strong  sulphuric  acid, 
and  heat  gently  on  a  sand  bath.  The  delivery  tube  should  pass  under 
water.  Carbon  monoxide  is  given  off  when  a  sufficiently  high  tempera- 
ture is  reached,  and  may  be  collected  in  bottles,  as  with  hydrogen.  Care 
must  be  taken  to  avoid  applying  too  much  heat,  as  the  reaction  would 
become  very  violent,  and  it  should  be  remembered  that  the  gas  must 
not  be  inhaled,  as  it  is  poisonous.  Collect  several  bottles  of  the  gas, 
and  test  them  as  follows:  To  one  add  lime-water;  noprecipitate  results. 
Touch  a  lighted  taper  to  another;  the  gas  burns  with  a  blue  flame. 
Pour  lime-water  into  the  bottle  in  which  this  combustion  took  place, 
and  notice  that  a  precipitate  now  forms,  showing  that  in  the  combustion 
carbon  dioxide  is  produced.  This  experiment  should  be  made  in  a  fume 
closet. 

The  reaction  between  the  ferrocyanide  and  sulphuric 
acid  is  somewhat  complex,  but  its  essential  features  may 
be  written  in  this  manner: 


K4Fe(CN)8+6H2S04 

6CO+2K2S04-f3(NH4)2S04+FeS04. 

The  six  molecules  of  water  represented  come  from  the 
crystallized  ferrocyanide  and  the  commercial  acid.  Carbon 
monoxide  is  made  very  readily  by  the  decomposition  of  ox- 
alic acid,  which  is  easily  brought  about  by  heating  with 
strong  sulphuric  acid  : 

H2C204  =  H20+CO+C02. 

The  sulphuric  acid  acts  merely  as  a  dehydrating  agent 
to  separate  the  elements  of  water.  The  evolved  gas  is 
passed  through  a  wash  bottle  containing  solution  of  caus- 
tic alkali  to  absorb  the  carbon  dioxide,  leaving  the  mon- 
oxide practically  pure. 

The  gas  is  easily  made  by  passing  a  current  of  the  di- 
oxide through  a  hot  porcelain  or  iron  tube  filled  with  bits 
of  broken  charcoal.  If  the  temperature  is  high  enough  a 
reduction  takes  place  in  this  manner  : 


The  heat  of  an  ordinary  combustion  furnace  is  sufficient 
for  the  reduction.  This  reaction  is  one  which  takes  place  in 
the  blast  furnace  when  iron  ores  are  smelted  and  also  in  an 
ordinary  coal  fire.  The  blue  flame  seen  above  the  coal  is 
that  of  the  burning  CO.  At  the  base  of  the  stove  or  fur- 


214  GENERAL  CHEMISTRY. 

nace  CO2  is  produced  to  be  decomposed  above  by  the  ex- 
cess of  carbon.  In  processes  of  incomplete  combustion, 
where  carbon  burns  in  a  limited  supply  of  air,  for  instance, 
the  gas  is  often  formed.  It  is  present  in  ordinary  illumin- 
ating gas  and  is  the  chief  substance  which  gives  to  this 
gas  poisonous  properties.  It  escapes  sometimes  from 
stoves  with  a  coal  fire,  when  the  exit  for  the  products  of 
combustion  is  insufficient. 

Properties.  As  shown  above,  the  gas  burns  readily 
with  oxygen.  It  is  but  slightly  soluble  in  water  and  does 
not  combine  with  alkali  solutions  as  the  dioxide  does.  It 
possesses  the  very  important  property  of  forming  a  pecu- 
liar stable  compound  with  the  haemoglobin  of  the  blood, 
which  greatly  interferes  with  the  oxygen  carrying  power  of 
the  latter,  and  to  this  its  marked  poisonous  action  is  due. 
The  gas  is  free  from  odor  when  pure;  it  has  recently  been 
condensed  to  a  liquid  by  application  of  pressure  at  a  low 
temperature.  One  liter  of  the  gas  under  normal  condi- 
tions weighs  1.2566  grams. 

ILLUMINATING  GAS. 

Brief  experiments  given  above  illustrate  the  produc- 
tion of  gases  by  the  dry  distillation  of  wood  and  coal.  On 
the  large  scale  both  substances  have  been  employed  in  the 
manufacture  of  gases  for  illuminating  purposes.  At  the 
present  time  two  general  processes  are  in  use  for  the  pro- 
duction of  gas  for  heating  and  lighting.  These  will  be 
briefly  described.  In  one  case  soft  coal  is  distilled  in  large 
retorts  and  this  process  is  now  known  as  the  "old  gas  proc- 
ess," while  in  the  other  case  steam  is  decomposed  by  hot 
carbon  under  certain  conditions,  and  this  in  various  modi- 
fications is  known  as  the  "new  gas  process." 

Old  Process.  The  details  of  this  are  best  explained 
with  reference  to  the  illustration,  Fig.  23. 

The  retorts,  of  which  several  are  built  in  a  "bank" 
over  one  furnace,  are  shown  at  C;  they  are  charged  with 
soft  coal  and  closed  by  luted  doors.  The  fire  on  the 


GENERAL  CHEMISTRY. 


215 


216  GENERAL  CHEMISTRY. 

grate,  A,  brings  the  coal  in  the  retorts  to  the  point  of 
decomposition.  The  hot  gases  and  vapors  pass  up  through 
vertical  pipes  connected  with  the  front  of  each  retort  and 
enter  a  larger  pipe  known  as  the  hydraulic  main.  This 
hydraulic  main,  B,  is  half  filled  with  water  through  which 
the  gases  must  bubble.  This  water  serves  several  pur- 
poses, one  of  which  is  to  prevent  ingress  of  air  tothe  system 
when  the  retorts  are  opened  from  time  to  time  to  receive 
fresh  charges  of  coal.  After  leaving  the  hydraulic  main  the 
gas  passes  through  a  series  of  pipes  known  as  condensers, 
where  it  is  cooled  considerably  and  loses  a  large  part  of 
the  tarry  matter  brought  over  from  the  retorts.  This  tar 
and  condensed  water  run  into  a  well  through  a  trap,  H. 
Then  the  gas  goes  through  a  coke  tower  called  the  scrub- 
ber, to  be  further  freed  from  tar  and  ammoniacal  salts  by 
means  of  water  trickling  over  the  coke.  Beyond  the  coke 
tower  the  gas  enters  another  purifying  appliance  which  is 
essentially  a  large  box  containing  a  number  of  trays  filled 
with  lime.  The  gas  in  passing  back  and  forth  over  this 
lime  gives  up  a  large  part  of  the  hydrogen  sulphide  and 
carbon  dioxide  contained  in  it  and  is  then  ready  for  burn- 
ing. From  the  lime  box  it  passes  by  the  pipe,  S,  to  the 
gas  holder,  G,  and  then  to  the  city  mains  by  the  pipe,  S'. 

The  student  will  understand  that  this  illustration  is 
merely  a  diagram,  showing  the  essential  steps  in  the  proc- 
ess only.  The  hydrocarbon  gases,  as  liberated  in  the  re- 
torts, are  mixed  with  sulphur  compounds  from  the  ferrous 
sulphide  in  all  soft  coal,  and  with  ammoniacal  salts  from 
the  nitrogen  and  oxygen  in  the  coal.  A  large  amount  of 
tar  is  always  formed,  and  these  several  impurities  must  be 
removed  before  the  gas  is  suitable  for  burning  in  houses. 
The  purification  of  the  product  is  therefore  the  most  diffi- 
cult part  of  the  process,  and  various  appliances  have  been 
introduced  for  the  purpose,  but  their  important  character- 
istics are  shown  above.  The  tar  which  condenses  in  dif- 
ferent parts  of  the  plant  is  saved  and  distilled.  Among 
the  valuable  products  obtained  in  the  distillation  are  ben- 
zene, toluene,  phenol,  naphthalene  and  anthracene,  which 
are  largely  used  at  the  present  time  in  the  manufacture  of 
compounds  valuable  in  medicine,  in  the  arts  as  dye-stuffs, 


GENERAL  CHEMISTRY. 


217 


and  elsewhere.  Most  of  our  ammonia  and  salts  of  am- 
monium are  obtained  from  the  wash  waters  of  the  gas 
works,  and  finally  the  coke  left  in  the  retorts  after  the  com- 
pletion of  the  process  is  employed  as  a  valuable  fuel.  The 
average  composition  of  the  purified  gas  as  made  from  soft 
coal  is  shown  by  the  following  table,  the  gas  being  consid- 
ered in  the  dry  condition  : 


1. 

2. 

3. 

4. 

Hydrogen  

46.2 

39.8 

51.3 

46  0 

34.8 

43.1 

36  5 

39  5 

Carbon  monoxide  

8.9 

4.7 

4  5 

5  0 

Heavy  hydrocarbons  
Carbon  dioxide 

6.1 
1  5 

4.8 
3  0 

4.9 
1  i 

5-9 
1  2 

Nitrogen 

2  1 

4  6 

1  4 

1  9 

Oxygen  

0.4 

0.3 

0.5 

For  many  years  all  the  illuminating  gas  made  in  this 
country  was  produced  essentially  by  this  method,  but  since 
1878  another  general  process  has  come  into  favor  by  which 
the  larger  part  of  the  gas  consumed  in  cities  is  now  made. 

The  New  Process.  While  many  kinds  of  plant  are 
employed  in  practice,  the  foundation  principle  in  all  the 
newer  methods  is  this:  At  a  high  temperature  steam  is 
decomposed  by  carbon  with  formation  of  hydrogen,  carbon 
monoxide  and  carbon  dioxide,  the  proportions  of  the  last 
two  depending  on  the  temperature  and  excess  of  steam. 
We  have  these  reactions: 


H2  +CO 

2H2O-f-C  =  2H3+CO2. 

In  practice  anthracite  coal  is  brought  to  a  white  heat  in 
a  retort  by  the  aid  of  a  blast  of  air.  When  this  stage  is 
reached  the  air  is  shut  off  and  live  steam  blown  in.  This 
acts  on  the  coal,  yielding  the  products  named.  The  gase- 
ous mixture  on  leaving  the  retort  passes  through  cooling 
and  washing  pipes  and  through  lime  boxes  for  the  absorp- 
tion of  the  carbon  dioxide.  What  is  left  is  nearly  pure 
hydrogen  and  carbon  monoxide  and  constitutes  what  is 


218  GENERAL  CHEMISTRY. 

known  as  pure  water  gas  or  fuel  gas.  It  burns  with  a  blue 
flame,  and  has  been  employed  for  heating  and  to  some 
extent  for  illumination  by  the  aid  of  a  special  burner  which 
becomes  incandescent  under  the  action  of  the  burning  gas. 
But  most  of  the  product  is  " carbonized  "  or  "enriched" 
as  made,  and  furnished  directly  as  an  illuminating  gas.  To 
accomplish  this  various  methods  have  been  patented.  They 
amount  essentially  to  this:  A  stream  of  light  petroleum  is 
injected  into  the  retort  in  which  the  first  reaction  takes 
place  or  into  aconnected  following  chamber,  called  a  "super- 
heater," through  which  the  hot  gases  pass.  In  either  case 
the  petroleum  is  decomposed  by  the  high  temperature, 
yielding  permanent,  light,  gaseous  products  very  similar 
to  those  found  in  the  soft  coal  gas.  Any  excess  of  the  oil 
not  decomposed  condenses  in  the  cooling  pipes  and  scrub- 
bers and  has  no  further  action.  In  its  composition  water 
gas  is  mainly  distinguished  from  soft  coal  gas  by  contain- 
ing less  methane  but  a  much  larger  amount  of  carbon  mon- 
.  oxide.  When  inhaled  it  is  therefore  more  rapidly  fatal, 
but  otherwise  in  many  respects  is  superior  to  the  soft  coal 
product. 

Of  the  many  minor  processes  employed  locally  in  pro- 
duction of  gases  for  lighting  no  mention  need  be  made  here, 
as  they  have  little  technical  importance. 

Gas  Burners.  Gas  is  used  for  two  essentially  different 
purposes;  first,  for  the  production  of  light,  and  secondly  for 
heating.  The  illuminating  burner  is  so  constructed  that 
air  is  mixed  with  the  gas  at  the  point  of  combustion  only. 
The  combustion  takes  place  from  the  outside  and  is  rela- 
tively slow.  A  part  of  the  carbon  is  thrown  into  the  re- 
duced or  free  condition,  and,  becoming  incandescent, 
makes  the  flame  luminous. 

In  burners  employed  in  heating,  air  is  drawn  into  the 
gas  stream  by  some  device  before  the  point  of  combustion 
is  reached.  This  air  mixes  thoroughly  with  the  gas  and 
with  the  aid  of  the  outside  air  provides  oxygen  enough  for 
perfect  and  rapid  combustion  of  the  carbon  as  well  as  hy- 
drogen present.  As  no  free  carbon  is  produced  here  we 
have  no  incandescence.  The  common  Bunsen  burner  was 


GENERAL  CHEMISTRY.  219 

the  first  and  is  still  the  best  illustration  of  this  class  of  de- 
vices. 

For  the  production  of  higher  temperatures  than  can  be 
secured  by  ordinary  combustion  the  blast  lamp  is  em- 
ployed. In  this  a  current  of  air  or  oxygen  is  blown  by  a 
bellows  or  otherwise  into  the  center  of  the  gaseous  stream 
so  as  to  hasten  the  oxidizing  action.  The  best  effect  is  ob- 
tained with  pure  hydrogen  and  oxygen,  but  oxygen  and 
coal  gas  are  very  commonly  used,  as  in  the  production  of 
the  calcium  light.  In  the  laboratory  blast  lamp  air  is 
blown  into  the  common  coal  gas. 

In  recent  years  another  form  of  lamp  has  become  com- 
mon, and  in  this  the  illumination  is  secured  by  the  incandes- 
cence of  a  gauze  cylinder  or  cone  made  of  the  oxides  of  sev- 
eral rare  metals.  A  flame  resembling  the  Bunsen  flame 
plays  against  this  gauze  cylinder,  bringing  it  up  to  the 
white  hot  condition  when  it  emits  a  very  brilliant  light. 
The  material  of  the  gauze  cylinder  consists  usually 
of  a  mixture  of  the  oxides  of  thorium,  lanthanum, 
cerium  and  other  rare  metals.  These  oxides  become  in- 
candescent at  a  relatively  low  heat.  Lime  becomes  incan- 
descent when  heated  in  the  oxyhydrogen  flame,  and  the  light 
so  produced  is  called  the  calcium  light,  lime  light  or 
Drummond  light. 

Safety  Lamp.  It  is  a  matter  of  common  observation 
in  the  laboratory  that  the  Bunsen  burner  flame  may  ap- 
pear above  wire  gauze  and  not  strike  through,  and  also 
that  when  burning  below  it  is  often  unable  to  pass  through 
and  burn  above.  The  explanation  of  this  is  found  in  the 
fact  that  the  gauze  is  so  perfect  a  conductor  of  heat  that 
the  temperature  of  the  combustible  gas  on  the  opposite 
side  of  the  gauze  from  the  flame  does  not  reach  the  kin- 
dling point.  This  is  easily  shown  by  lamp  flames  as  well  as 
with  those  from  gas.  The  principle  was  recognized  and  ex- 
plained by  Humphrey  Davy  and  applied  by  him  in  the  con- 
struction of  the  safety  lamp.  This  is  an  oil  lamp  with  a 
chimney  of  fine  iron  gauze.  The  top  and  bottom  of  the 
chimney  are  not  open,  but  made  of  gauze  so  that  the  flame  is 
surrounded.  When  the  lighted  lamp  is  taken  into  a  mine 


220  GENERAL  CHEMISTRY. 

containing  marsh  gas  this  passes  with  the  air  through  the 
meshes  and  causes  a  slight  explosion.  But  the  conducting 
power  of  the  wire  keeps  the  temperature  of  the  burning 
gas  below  the  point  at  which  the  explosive  mixture  out- 
side could  be  ignited,  giving  the  miner  time  to  seek  a  place 
of  safety.  The  lamp  gives  a  warning  of  the  presence  of 
marsh  gas;  if  kept  burning  long  enough  in  such  an  atmos- 
phere the  gauze  chimney  may  in  time  become  hot  enough 
to  communicate  the  combustion  to  the  outside. 

CARBON  AND  HYDROGEN. 

Many  hundreds  of  compounds  of  these  elements  are 
known,  some  of  them  being  natural  products,  while  others 
are  made  by  laboratory  operations.  They  are  called  hydro- 
carbons. Ordinary  crude  petroleum  consists  essentially  of 
bodies  of  this  class,  while  other  natural  substances  contain 
them  also.  The  part  of  chemistry  known  as  organic  chem- 
istry is  concerned  with  a  study  of  these  substances  and 
others  related  to  or  derived  from  them.  A  few  will  be 
studied  here  as  illustrations. 

Methane  or  Marsh  Gas. 

This  is  a  light,  gaseous  body,  containing  75  per  cent  of 
carbon  and  25  per  cent  of  hydrogen  by  weight.  It  is  found 
in  the  gases  escaping  from  oil  wells  and  makes  up  a  large 
proportion,  often  95  per  cent,  of  the  so-called  natural  gas 
found  in  many  parts  of  the  country.  It  is  often  given  off 
from  marshy  ground  or  stagnant  pools,  hence  the  name, 
and  is  formed  there  by  the  decay  of  organic  matter.  In 
the  laboratory  we  can  make  it  by  the  following  method: 

Ex.  144.  Prepare  a  mixture  of  1  part  of  thoroughly  dried  (dehy- 
drated) sodium  acetate  and  three  parts  of  soda-lime.  Charge  about  20 
Gm.  of  this  into  a  gas-pipe  retort  arranged  as  in  the  experiments  on  the 
production  of  coal  and  wood  gases.  When  a  strong  heat  is  applied  to 
the  retort  a  decomposition  of  the  mixture  takes  place  and  methane  is 
given  off.  This  is  collected  over  water.  Fill  one  or  two  bottles  and 
test.  The  gas  burns  with  a  bluish  flame  if  pure,  but  as  here  made  has 
often  a  yellowish  color  from  the  presence  of  sodium  compounds  carried 
over  in  traces  from  the  mixture  in  the  retort.  The  gas  is  very  light  and 


GENERAL  CHEMISTRY.  221 

can  be  poured  upward,  as  was  hydrogen.    Test  with  lime-water  for  car- 
bon dioxide  among  the  products  of  combustion,  when  the  gas  is  burned. 

Methane  is  a  constituent  of  illuminating  gas,  amount- 
ing often  to  40  per  cent  of  the  whole  by  volume.  It  col- 
lects sometimes  in  deep  coal  mines  and  is  there  known  as 
fire  damp,  giving  rise  frequently  to  violent  explosions. 

The  reaction  which  takes  place  when  methane  is  made 
by  the  above  process  is  this: 


CH4 

Sodium  i_  Sodium    —     Sodium      _i_  Mpthan^ 

acetate          "Thydroxide  —   carbonate  ~T  r 

The  sodium  hydroxide  is  contained  in  the  soda-lime. 
When  methane  burns,  water  and  carbon  dioxide  are  formed 
as  here  illustrated: 

CH4+202=C02-f-2H20. 

One  volume  of  methane  requires  two  volumes  of  oxy- 
gen for  complete  combustion,  or  ten  volumes  of  air.  The 
gas  is  insoluble  in  water,  practically,  and  is  not  absorbed 
by  alkali  solutions. 

Ethylene. 

This  a  colorless  gas  which  is  formed  in  small  quantity 
in  illuminating  gases.  In  the  laboratory  it  is  most  readily 
made  by  heating  a  mixture  of  alcohol  and  strong  sul- 
phuric acid. 

Ex.  145.  Pour  about  20  Cc.  of  alcohol  into  a  flask  of  300  to  400 
Cc.  capacity  and  add,  a  little  at  a  time,  and  with  agitation  of  the  flask, 
about  80  to  100  Cc.  of  strong  sulphuric  acid.  Close  the  flask  with  a 
stopper  carrying  a  safety  tube  and  a  delivery  tube.  Lead  the  latter  into 
a  Woulfe  bottle  as  in  Ex.  68  on  the  production  of  HC1.  In  the  present 
case  half  fill  the  first  bottle  with  sulphuric  acid,  into  which  the  delivery 
tub*  dips,  and  the  second  bottle  with  a  solution  of  sodium  hydroxide. 
Now  heat  the  flask  on  a  sand-bath  very  gradually  until  the  evolution  of 
gas  begins,  as  shown  by  the  rapid  escape  of  bubbles  through  the  alkali 
solution.  Then  regulate  the  flame  so  as  to  prevent  too  rapid  a  reaction 
in  the  flask,  as  this  would  lead  to  frothing  of  the  liquid,  and  overflow 
through  the  funnel  tube  and  delivery  tube-  Remove  the  lamp  entirely 
if  necessary.  As  generated,  the  gas  is  not  pure  but  on  passing  through 
the  acid  of  the  first  bottle  and  the  alkali  of  the  second  it  gives  up  most 
of  the  contaminations,  consisting  mainly  of  carbon  dioxide,  sulphurous 


222  GENERAL  CHEMISTRY. 

oxide  and  ether  vapor.  From  the  second  Woulfe  bottle  a  delivery  tube 
leads  to  a  trough  holding  inverted  bottles  full  of  warm  water  for  the  col- 
lection of  the  gas.  Fill  two  or  three  bottles  and  test  the  gas  by  burning  it. 
After  lighting  the  gas  at  the  mouth  of  a  bottle  pour  water  through  the 
flame  to  force  the  gas  up  to  the  air  more  rapidly.  A  bright  illuminating 
flame  is  produced.  The  gas  may  also  be  burned  at  the  end  of  the  deliv- 
ery tube,  or  after  passing  through  a  common  lava  gas  tip. 

This  gas,  although  being  present  in  but  small  propor- 
tion in  common  illuminating  gas,  is  important  because  of 
its  light  giving  property.  Hydrogen,  methane  and  carbon 
monoxide  make  up  about  85  to  90  per  cent  of  the  volume  of 
our  common  gas,  but  burn  with  nonluminous  flames. 
Ethylene  and  several  other  gaseous  substances  present  in 
small  amount  burn  with  highly  luminous  flames  and  are 
valuable  on  that  account. 

It  has  been  shown  that  sulphuric  acid  has  a  marked 
affinity  for  water  and  that  it  can  even  decompose  organic 
substances  to  combine  with  it.  The  decomposition  of 
alcohol  is  an  illustration  of  this.  Alcohol  is  a  compound 
of  carbon,  hydrogen  and  oxygen  having  the  formula 
C2H6O.  The  strong  acid  decomposes  this,  taking  out 
H2O  and  leaving  C2H4  : 

C2H60— H20  =  C2H4. 

Ethylene. 

Acetylene. 

This  is  another  interesting  compound  of  carbon  and 
hydrogen,  with  the  composition  expressed  by  the  formula 
C2H2.  It  is  found  in  illuminating  gas  and  is  made  in 
processes  of  imperfect  combustion  of  some  other  gases  or 
vapors.  Recently  it  has  received  much  attention  from 
chemists  because  of  the  fact  that  by  a  new  method  it  can 
be  produced  at  small  cost  and  in  any  desired  quantity. 
As  it  is  a  highly  valuable  illuminant  this  discovery  is*of 
importance.  It  is  made  in  this  new  process  by  the  decom- 
position of  a  compound  known  as  calcium  carbide,  by 
means  of  water.  An  experiment  will  show  this. 

Ex.  146.  Arrange  a  bottle  exactly  as  for  the  generation  of  hydro- 
gen. The  bottle  should  have  a  capacity  of  about  150  Cc.  Put  15  to  20 
Gm.  of  the  calcium  carbide  in  the  bottle  and  add  strong  alcohol  to  make 


GENERAL  CHEMISTRY.  223 

a  layer  about  1  Cm.  in  depth.  Close  the  bottle  with  the  cork,  holding 
a  funnel  tube  and  a  delivery  tube,  the  latter  leading  to  a  trough  with 
inverted  collecting  bottles.  The  lower  end  of  the  funnel  tube  should 
dip  below  the  surface  of  the  alcohol.  Now  pour  in  a  little  water  slowly 
and  continue  the  addition  until  a  rapid  evolution  of  gas  begins.  Allow 
the  first  portions  to  escape  into  the  air  and  then  collect  some  for  tests. 
Notice  that  it  burns  with  a  heavy,  smoky  flame.  If  the  gas  is  burned 
through  a  lava  gas  tip,  it  is  bright  and  luminous.  Lead  a  little  of  the 
gas  from  the  generating  bottle  into  a  flask  containing  a  solution  of 
cuprous  chloride  in  ammonia  water.  A  characteristic  red  precipitate  is 
formed. 

The  gas  has  a  peculiar  disagreeable  odor  r.nd  is  poison- 
ous when  inhaled  in  quantity.  The  reaction  by  which  it 
was  formed  is  illustrated  as  follows  : 

CaC2+2H2O:=CaO2H2-fC2H2 

«*?£  +  Wa<er    =   hydride    +Acetylene. 

Acetylene  may  be  condensed  to  the  liquid  condition 
without  practical  difficulty  and  to  some  extent  is  used  in 
that  form.  It  is  preferable  and  safer,  however,  to  burn  it 
as  generated,  at  the  ordinary  pressure.  The  carbide  from 
which  it  is  made  is  a  good  illustration  of  a  class  of  sub- 
stances which  have  recently  been  made  in  quantity  by  aid 
of  the  high  heat  of  the  electric  arc.  In  this  case  a  mixture 
of  powdered  coke  and  lime  is  exposed  to  the  heat  of  the 
arc  with  the  formation  of  carbide  and  carbon  monoxide. 


CARBON,  HYDROGEN  AND  OXYGEN. 

These  compounds,  like  those  of  carbon  and  hydrogen 
alone,  are  very  numerous.  The  starches,  sugars,  gums, 
resins,  alcohols,  fats  and  many  other  common  substances 
contain  these  three  elements.  At  one  time  it  was  sup- 
posed that  these  bodies  could  be  produced  only  by  the 
agency  of  living  plants  or  animals.  We  know  now  that 
this  is  not  true,  as  many  combinations  of  these  three 
elements  are  easily  made  by  laboratory  operations.  A  few 
simple  illustrations  will  be  given  here  of  the  preparation 
and  properties  of  some  so  called  organic  substances. 


224  GENERAL  CHEMISTRY. 

Starch  will  be  taken  as  the  starting  point,  as  it  is  a  com- 
mon substance  and  is  used  for  many  purposes. 

Ex.  147.  Rub  about  10  grams  of  starch  to  a  thin  cream  with  a 
little  water.  Then  add  water  to  make  a  volume  of  about  200  Cc.,  and 
boil  the  mixture  to  form  a  paste.  Add  now  5  Cc.  of  dilute  sulphuric 
acid  and  boil  three  hours,  or  more,  in  a  flask.  Add  a  few  drops  of  hot 
water,  from  time  to  time,  to  compensate  for  that  lost  by  evaporation. 
At  the  end  of  the  boiling  it  will  be  noticed  that  the  mixture  is  clear,  the 
starch  having  been  dissolved.  Next  add  to  the  hot  liquid  some  fine 
marble  dust  and  shake  the  mixture  thoroughly.  Continue  the  addition 
as  long  as  carbon  dioxide  is  given  off,  by  the  action  of  the  free  sulphuric 
acid  on  the  calcium  carbonate.  The  object  of  adding  the  marble  dust 
is  to  neutralize  the  acid.  Allow  the  mixture  to  cool,  dilute  it  with  some 
water,  and  filter  it  into  a  clean  beaker.  If  the  filtrate  has  an  acid  taste 
add  a  little  more  marble  dust  to  it,  warm,  allow  to  settle,  and  filter 
again.  Evaporate  this  filtrate  to  dryness,  best  over  a  water-bath.  At 
any  rate  the  temperature  must  not  get  high  enough  to  scorch  the  sub- 
stance as  it  becomes  concentrated.  From  time  to  time  take  up  a  drop 
of  the  liquid  on  a  glass  rod,  cool  and  taste  it.  By  the  action  of  the  sul- 
phuric acid  on  the  starch  a  sugar,  known  as  dextrose,  is  produced.  The 
acid  itself  is  but  slightly  altered  in  the  process;  most  of  it  remains  until 
the  end  of  the  boiling,  and  then  it  is  combined  with  the  marble  dust  to 
form  insoluble  calcium  sulphate,  which  is  removed  in  the  filtration.  The 
weak  sugar  solution  is  concentrated  until  its  nature  becomes  apparent 
by  the  taste. 

Starch  is  represented  by  the  formula  C6H10O5,  and  the 
dextrose  by  the  formula  C6H12O6.  The  weak  acid  has 
apparently  added  H2O  to  the  starch,  which  is  a  behavior 
quite  the  reverse  of  that  of  the  strong  acid,  already  shown. 
The  reaction  illustrated  here  is  identical  with  one  followed 
on  the  large  scale,  in  which  commercial  glucose  is  made  by 
heating  starch  paste  with  weak  sulphuric  acid,  in  closed 
vessels  under  steam  pressure. 

To  show  some  of  the  properties  of  the  sugar  formed  in 
the  experiment  dissolve  it  in  water  and  test  the  solution. 

Sugar  Tests.  We  have  several  very  sharp  reactions 
by  which  we  are  able  to  recognize  dextrose  in  a  solution. 
One  of  the  best  of  these  will  be  given  here. 

Ex.  148.  Add  to  a  few  Cc.  of  the  sugar  solution  an  equal  volume  of 
strong  potassium  hydroxide  solution,  and  then  a  few  drops  of  a  weak 
solution  of  copper  sulphate.  A  blue  color  appears,  much  deeper  than 
that  of  the  copper  sulphate.  Now  boil  the  liquid.  On  heating,  it 
becomes  turbid  and  greenish  yellow.  A  precipitate  forms  which  darkens 


GENERAL  CHEMISTRY.  225 

and  becomes  heavier  until  at  the  boiling  temperature  it  is  bright  red. 
This  precipitate  consists  of  a  combination  of  copper  with  oxygen  known 
as  cuprous  oxide,  and  represented  by  the  formula  Cu2O.  The  sugar 
acts  here  as  a  reducing  substance,  but  the  nature  of  the  reaction  cannot 
be  explained  at  this  point.  The  test,  as  carried  out  here,  is  known  as 
Trommels  test.  A  very  similar  behavior  is  shown  on  boiling  the  sugar 
solution  with  Fehling's  solution,  which  also  contains  copper.  The  prep- 
aration of  this  solution  is  given  in  books  on  analytical  chemistry.  It 
is  essentially  a  solution  of  copper  hydroxide  with  an  alkali  and  a  tar- 
trate. 

These  tests  are  characteristic  of  dextrose,  and  some 
other  sugars,  but  not  of  cane  sugar.  When  a  solution  of 
cane  sugar  is  boiled  with  a  weak  acid,  however,  it  is  con- 
verted into  dextrose  and  a  similar  substance  known  as 
levulose  and  this  mixture  responds  to  the  test.  Prove  this 
by  an  experiment. 

Ex.  149.  Dissolve  about  5  Gm.  of  pure  cane  sugar  in  water,  in  a 
clean  test-tube.  Divide  the  solution  into  two  parts.  Apply  the  Trom- 
mer  or  Fehling  test  to  one  of  these  and  observe  that  no  red  precipitate  is 
formed.  Then  boil  the  other  half  of  the  solution  with  a  few  drops  of 
strong  hydrochloric  acid  some  minutes,  add  some  alkali  to  neutralize 
the  acid, 'and  apply  the  test  as  given.  A  bright  red  precipitate  appears 
almost  immediately. 

These  tests  are  of  the  highest  importance  and  are  em- 
ployed frequently  in  analytical  chemistry. 

From  starch  we  made  dextrose  by  action  of  a  weak  acid 
and  heat.  It  will  next  be  shown  how  the  sugar  may  be 
converted  into  something  else  of  equally  great  importance. 
But  for  this  we  will  need  a  greater  quantity  of  sugar  than 
was  made  in  the  above  experiment.  Common  glucose 
syrup  or  molasses  may  be  used  for  the  purpose. 

Fermentation.  When  a  solution  of  molasses  or  glu- 
cose, not  too  concentrated,  is  mixed  with  a  little  yeast,  or 
exposed  some  time  to  the  air,  it  undergoes  a  change  which 
we  call  fermentation.  The  solution  becomes  lighter,  gives 
off  a  gas  which  we  recognize  as  carbon  dioxide,  and  emits 
the  characteristic  odor  of  alcohol.  This  change  is  pro- 
duced by  the  agency  of  what  is  termed  a  ferment.  There 
are  several  kinds  of  ferments  ;  but  the  one  concerned  here, 
the  yeast  ferment,  consists  of  minute  vegetable  cells  which 


226  GENERAL  CHEMISTRY. 

can  be  seen  under  the  microscope.  The  function  of  these 
cells  is  to  consume  sugar  and  produce  alcohol  and  carbon 
dioxide.  The  change  will  be  shown  by  an  experiment. 

Ex.  150.  Mix  100  Cc.  of  common  glucose  molasses  with  800  to 
900  Cc.  of  water,  and  with  the  mixture  nearly  fill  a  bottle  or  flask. 
Add  5  Cc.  of  brewer's  yeast,  or  some  compressed  yeast  previously 
soaked  in  water.  Close  the  bottle  with  a  perforated  stopper  and  con- 
nect it  with  a  Woulfe  bottle  and  soda-lime  tube  as  shown  in  the  next 
figure. 

Half  fill  the  Woulfe  bottle  with  lime-water.  Leave  the  apparatus 
two  or  three  days  in  a  moderately  warm  place.  Active  fermentation 
soon  begins  and  bubbles  pass  over  from  the  generating  bottle  into  the 
lime-water,  where  a  precipitate  is  soon  formed.  The  alkali  in  the  soda- 
lime  tube  serves  to  protect  the  lime-water  from  the  carbon  dioxide  of 


FIG.  24. 

the  air  at  the  beginning  of  the  experiment.  The  precipitate  which 
forms  at  first  in  the  lime-water  clears  up  later,  as  an  excess  of  the  gas 
passes  through.  At  the  end  of  three  days  disconnect  the  apparatus  and 
use  the  liquid  in  the  fermentation  bottle,  which  is  now  a  weak  solution 
of  alcohol,  for  the  following  experiments  : 

Ex.  151.  Divide  the  fermented  liquid  into  two  halves.  Allow  the 
one  half  to  stand  some  days  in  an  open  bottle,  exposed  to  the  air. 
With  the  other  half  proceed  as  follows  :  Fill  it  into  a  flask,  as  shown  on 
the  left  in  the  figure  below,  and  connect  this  with  a  second  flask  on  a 
water-bath  or  dish  of  water.  A  delivery  tube  from  this  second  flask 
leads  in  turn  to  a  small,  open  flask,  which  may  be  kept  cool  by  standing 
it  in  cold  water. 

The  first  flask  is  heated  on  a  sand-bath,  so  that  the  liquid  in  it  just 
boils.  The  alcohol  present  distills  with  water  over  into  the  second 


GENERAL  CHEMISTRY. 


227 


flask.  This  is  heated  to  a  lower  temperature,  in  order  to  distill,  as  far 
as  possible,  only  the  alcohol.  As  there  is  a  difference  of  20°  C.  in  the 
boiling  points  of  alcohol  and  water,  a  much  stronger  product  may  be 
driven  over  from  the  second  flask  than  from  the  first.  The  distilled  al- 
cohol with  water  collects  in  the  small  flask,  and  may  be  tested  by  burn- 
ing. At  the  end  of  the  distillation,  that  is,  after  about  a  third  of  the 
contents  of  the  first  flask  has  been  boiled  over,  pour  some  of  the  product 
from  the  small  flask  into  a  porcelain  dish  and  apply  a  light.  It  should 
burn  with  a  colorless  flame. 

On  the  large  scale  alcohol  is  produced  by  an  analogous 
process,  and  is  concentrated   by    distillation   until  it  has  a 


FIG.  25. 


strength  of  about  93  per  cent  by  volume.  The  proportions 
of  alcohol  and  carbon  dioxide  obtained  from  the  sugar  are 
represented  by  this  equation: 

C6H1206=2C2H60+2CO: 


Dextrose       — 


Alcohol  is  employed  for  many  purposes  in  the  labora- 
tory and  in  the  arts.  It  is  used  as  a  solvent  in  making  var- 
nishes, perfumes,  etc.,  and  also  as  the  starting  point  in  the 


228  GENERAL  CHEMISTRY. 

manufacture  of  many  other  substances.    Ether  and  chloro- 
form are  illustrations  of  bodies  made  from  alcohol. 

Acid  Fermentation.  We  have  yet  the  other  half  of 
our  alcoholic  liquid  to  examine.  Allow  it  to  stand  in  the 
air,  as  directed,  until  the  odor  of  alcohol  disappears  and  a 
sour  odor  takes  its  place.  Under  certain  circumstances, 
on  standing,  alcohol  becomes  converted  into  acetic  acid. 
The  change  is  much  more  rapid  if  to  it  is  added  a  little  of 
the  substance  known  as  mother  of  vinegar,  which,  like  the 
yeast,  is  a  ferment.  The  alcohol  ferment  and  the  acetic 
acid  ferment  are  present  in  nearly  all  atmospheres,  and 
bring  about  the  production  of  alcohol,  and  then  the  change 
into  vinegar,  in  fruit  juices  exposed  to  the  air. 

Ex.  152.  Treat  the  alcoholic  liquid,  which  has  become  sour,  in  this 
manner:  Filter  it  first  and  to  the  clear  filtrate  add  a  few  grams  of  mar- 
ble dust.  A  slow  solution  of  the  solid  will  be  observed,  as  the  acid 
liquid  forms  with  it  soluble  calcium  acetate,  with  escape  of  carbon  diox- 
ide. The  solution  is  aided  by  heat.  When  no  more  marble  dust  dis- 
solves filter  the  mixture  and  concentrate  the  filtrate  on  a  water-bath,  or 
sand-bath,  slowly.  On  a  sand-bath,  care  must  be  taken  not  to  char  the 
concentrated  product.  When  cold  the  nature  of  this  may  be  readily 
recognized  by  adding  a  little  strong  sulphuric  acid.  This  decomposes 
the  calcium  acetate,  as  it  does  other  salts  (chlorides,  nitrates,  carbon- 
ates, etc.),  with  liberation  of  acetic  acid,  which  is  recognized  by  the 
odor.  The  object  of  converting  the  acid  in  the  sour  liquid  into  calcium 
acetate  was  to  secure  a  product  which  could  be  concentrated  without 
loss  and  then  be  recognized.  In  the  original  liquid  the  acid  was  in  a 
highly  diluted  condition,  and  not  easily  recognizable. 

The  composition  of  acetic  acid  is  C2H4O2,  while  that 
of  the  alcohol  from  which  it  was  derived  is  C2H6O.  The 
vinegar  ferment,  therefore,  adds  oxygen  and  abstracts  hy- 
drogen, but  in  what  manner  this  is  done  we  cannot  ex- 
plain. Acetic  acid  is  an  important  substance,  occurring 
to  the  extent  of  3  or  4  per  cent  in  ordinary  vinegar. 
Vinegar  and  weak  acetic  acid,  for  certain  purposes,  are 
made  by  fermentation,  but  a  great  deal  of  our  strong  acid 
is  obtained  from  the  acid  liquid  produced  in  the  distillation 
of  wood,  and  known  in  crude  form  as  pyroligneous 
acid.  Acetic  acid  forms  a  class  of  salts  known  as 
acetates,  of  which  we  have  had  several  illustrations. 


GENERAL  CHEMISTRY.  229 

CARBON  AND  NITROGEN. 

Carbon  and  nitrogen  enter  into  important  combina- 
tions, one  of  which  will  be  briefly  referred  to.  This  is 
cyanogen  and  is  represented  by  the  symbols  CN.  This 
cyanogen  combines  with  metals,  forming  compounds 
known  as  cyanides.  Potassium  cyanide  is  an  illustration. 
This  has  the  composition  KCN.  In  this  group  the  car- 
bon and  nitrogen  seem  to  have  the  effect  of  a  single  ele- 
ment, for  instance,  chlorine.  Potassium  cyanide  is  in 
some  respects  analogous  to  potassium  chloride.  It  goes 
into  combinations  with  other  cyanides  very  readily,  yield- 
ing so-called  double  cyanides.  The  potassium-iron  cya- 
nides will  be  referred  to  later,  as  they  are  interesting  and 
important  compounds. 

A  solution  of  potassium  cyanide  is  used  in  photography 
as  a  fixing  agent  because  of  its  marked  solvent  action  on 
silver  chloride  or  bromide.  The  ordinary  plating  baths 
employed  in  gold  and  silver  plating  are  usually  made  of 
the  double  cyanides  of  potassium  and  silver  and  potas- 
sium and  gold  dissolved  with  an  excess  of  potassium  cya- 
nide. These  solutions  are  intensely  poisonous, 

As  a  chloride  is  decomposed  by  sulphuric  acid,  yield- 
ing hydrochloric  acid,  so  a  cyanide  may  be  decomposed, 
yielding  hydrocyanic  acid,  HCN,  as  illustrated  by  this 
equation: 


Hydrocyanic  acid  is  commonly  known  as  prussic  acid. 
It  is  a  very  volatile  liquid,  extremely  poisonous,  and  may 
be  handled  with  safety  only  in  dilute  solutions.  It  dissolves 
in  water  in  all  proportions.  A  fuller  discussion  of  the 
cyanogen  compounds  belongs  in  the  field  of  organic  chem- 
istry. 

CARBON  AND  CHLORINE. 

Carbon  forms  a  number  of  important  compounds  with 
chlorine,  or  with  chlorine  and  hydrogen,  which  are 
described  in  organic  chemistry.  The  common  substance 
known  as  chloroform  has  the  formula  CHC13.  It  is  some- 


230  GENERAL  CHEMISTRY. 

times  called  trichlormethane  and  may  be  looked  upon  as 
marsh  gas,  or  methane,  in  which  three  atoms  of  hydrogen 
are  replaced  by  three  chlorine  atoms.  Another  compound 
has  the  formula  CC14,  and  is  called  tetrachlormethane,  or 
tetrachloride  of  carbon.  Both  are  volatile  liquids  with  a 
characteristic,  pleasant  odor  and  are  employed  in  medicine 
as  anaesthetics. 

CARBON  AND  SULPHUR. 

Carbon  forms  a  very  important  compound  with  sulphur, 
known  as  carbon  disulphide,  CS2.  It  is  produced  by  pass- 
ing the  vapor  of  sulphur  over  hot  carbon  in  a  retort.  At 
the  same  time  several  other  compounds  are  formed,  some 
of  which  remain  with  the  disulphide  as  contaminations 
hard  to  separate.  Pure  CS2  is  a  liquid  with  a  pleasant 
ethereal  odor,  boiling  at  47°  and  having  a  density  of  1.29. 
It  is  but  slightly  soluble  in  water,  and  is  itself  an  excellent 
solvent  for  sulphur,  caoutchouc,  fats  and  other  substances. 
Because  of  this  behavior  it  has  many  uses  in  the  arts  and 
in  analytical  chemistry.  The  common  commercial  disul- 
phide has  a  very  disagreeable  odor  due  to  the  presence  of 
the  impurities  formed  in  the  process  of  manufacture.  The 
vapor  of  carbon  disulphide  is  very  inflammable  and  is  also 
poisonous  when  inhaled.  Care  must  therefore  be  observed 
in  handling  it. 

In  its  chemical  behavior  CS2  bears  some  relation  to 
CO2.  With  the  latter  this  reaction  occurs: 

Na2O-{-CO2=Na2CO3, 

and  an  analogous  combination  is  known  for  the  disul- 
phide. Sulphide  of  sodium  unites  with  it  in  this  man- 
ner: 

Na2S-fCS2=Na2CS3. 

The  compound,  Na2CS3,  is  called  sodium  thiocarbon- 
ate,  or  sometimes  sulphocarbonate.  It  may  be  decom- 
posed by  hydrochloric  acid,  yielding  an  acid,  H2CS3, 
corresponding  to  the  real  carbonic  acid,  and  called  thio- 
carbonic  acid. 


CHAPTER  XI. 


ATOMIC  AND  MOLECULAR  WEIGHTS. 
DALTON'S  WEIGHTS. 

IN  Chapter  III  a  brief  outline  of  the  atomic  theory  as 
suggested  by  Dalton  was  given.  It  is  intended  in  the 
present  chapter  to  go  a  little  more  fully  into  details  and 
explain  some  of  the  steps  by  which  chemists  have  passed 
from  the  views  of  the  earlier  writers  to  those  held  at  the 
present  time. 

It  was  Dalton's  idea  that  where  atoms  combine  to  form 
compounds,  such  combinations  must  be  in  the  simplest 
possible  proportions,  an  atom  of  one  with  an  atom  of  the 
other  in  most  cases.  Giving  hydrogen  unit  atomic 
weight,  it  would  follow,  therefore,  that  atomic  weights 
would  be  found  by  determining  the  weights  of  different 
bodies  which  unite  with  one  part  of  hydrogen. 

Let  M  represent  an  atom  of  one  substance  and  N  an 
atom  of  another;  then  a  combination  between  the  two 
would  be  in  the  proportion: 


with  x  andjy  generally  1.  The  atomic  weights  which  Dal- 
ton actually  found  were  inaccurate  because  of  his  faulty 
analytical  methods,  but  his  principle  would  lead  to  the  fol- 
lowing: 

Water  =HO  H  :  O::l  :  8.00.  O  =  8.00. 
Ammonia=HN  H  :  N::l  :  4.67.  N  =  4.67. 
Ethylene^HC  H  :  C::l  :  6.00.  C=6.00. 

If  8  parts  of  oxygen  unite  with  1  part  of  hydrogen  and 
4.67  parts  of  nitrogen  unite  with  the  same  weight  of  hydro- 


232  GENERAL  CHEMISTRY. 

gen  then  it  might  be  expected  that  when  oxygen  and  nitro- 
gen unite  with  each  other  the  proportion,  8  :  4.67,  would  rep- 
resent the  relation  of  their  weights.  But  this  is  not  the 
case  and  the  atomic  weights  as  given  appear  inconsistent; 
it  will  be  later  seen  that  this  is  due  to  a  false  assumption 
regarding  x  and  y  in  the  proportion,  xM  :  ^N,  above.  The 
first  atomic  weights  published  by  the  followers  of  Dalton 
suffered  many  arbitrary  corrections  with  the  hope  of  mak- 
ing them  consistent  among  themselves. 

COMBINATION  OF  GAS  VOLUMES. 

Several  chemists  had  undertaken  to  determine  the  pro- 
portions in  which  various  gases  unite  with  each  other,  but 
it  remained  for  Gay  Lussac  to  discover  the  extreme  sim- 
plicity of  these  relations.  It  has  been  shown  in  an  earlier 
chapter  that  2  volumes  of  hydrogen  unite  with  1  volume  of 
oxygen  to  form  2  volumes  of  water  vapor.  Through  the 
labors  of  Gay  Lussac  and  others  a  number  of  similar  simple 
relations  were  established  which  are  shown  in  the  follow- 
ing short  table: 

2  vols.  of  hydrogen  +1  vol.  of  oxygen   =2  vols.  of  water  vapor. 

3  vols.  of  hydrogen  +1  vol.  of  nitrogen=2  vols.  of  ammonia. 

2  vols.  of  sulphurous  oxide-f  1  vol.  of  oxygen    =2  vols.  of  sulphuric  oxide. 
2  vols.  of  nitrogen  +1  vol.  of  oxygen   =2  vols.  of  nitrous  oxide. 

Somewhat  later  other  very  simple  relations  were  found, 
among  them  these: 

i  vol.   of  hydrogen  -j-  i  vol.  of  chlorine  =  2  vols.  of  hydrochloric   acid. 

4  vols.  of  hydrogen  -f  i  vol.  of  carbon  (vapor)  =  2  vols.  of  marsh  gas. 

Gay  Lussac  pointed  out  the  fact  that  in  all  cases  inves- 
tigated the  volume  of  product  formed  bears  a  very  simple 
relation  to  the  sum  of  the  components.  This  is  shown  in 
the  examples  quoted. 

The  question  was  naturally  asked:  What  is  the  reason 
for  this  combination  of  gases  in  simple  volume  propor- 
tions? Answers  of  the  highest  importance  in  the  history 
of  chemistry  were  given  in  two  directions.  The  first  to 
correctly  interpret  the  gas  experiments  of  Gay  Lussac, 
Humboldt  and  Davy  was  Amadeo  Avogadro,  a  professor  in 


GENERAL  CHEMISTRY.  233 

the  University  of  Turin.  In  1811  he  published  an  article 
in  which  he  deduced  from  these  experiments  the  law  that 
equal  volumes  of  gases,  under  like  conditions,  contain  the  same 
number  of  particles  or  integrant  molecules.  Unfortunately 
the  value  of  the  observation  of  Avogadro  was  not  recog- 
nized for  many  years.  More  will  be  said  about  it  pres- 
ently. 

THE  VOLUME  THEORY  OF  BERZELIUS. 

Among  the  ablest  chemical  investigators  working  in  the 
earlier  years  of  the  century  there  must  be  mentioned  Ber- 
zelius.  At  the  time  when  Dalton  announced  his  doctrine 
of  combinations  through  atoms  of  constant  weight  Berze- 
lius  was  already  engaged  in  the  analysis  of  chemical  com- 
pounds, and,  accepting  Dalton's  idea,  he  immediately 
began  the  task  of  fixing  atomic  weights  by  determining  the 
proportions  in  which  various  substances  combined  with 
oxygen  and  hydrogen.  Dalton  had  suggested  hydrogen  as 
the  basis  of  atomic  weights,  but,  inasmuch  as  the  known 
oxygen  compounds  were  more  numerous  than  those  of  hy- 
drogen and  more  readily  examined,  he  proposed  the  oxy- 
gen atom  as  the  unit  and  placed  its  weight  arbitrarily  at 
100.  In  addition  to  this  he  announced  two  important  prin- 
ciples to  be  employed  in  fixing  atomic  weights.  One  of 
these  he  derived  from  the  discoveries  of  Gay  Lussac  and 
his  followers  just  mentioned.  Berzelius,  like  Avogadro, 
recognized  immediately  the  great  importance  of  these  sim- 
ple volume  relations  in  the  combination  of  gases.  This 
was  his  interpretation  of  them.  To  unite  in  this  manner, 
he  reasoned,  equal  gas  volumes  must  contain  the  same  number 
of  atoms  under  like  conditions  of  temperature  and  pressure. 
From  this  it  would  follow  that  the  atomic  weights  of  the 
gaseous  substances  must  be  related  to  each  other  as  are 
their  specific  gravities.  This  he  called  his  volume  theory 
in  atomic  weight  determination.  It  will  be  observed  that 
he  made  no  distinction  between  elementary  and  compound 
gases,  a  distinction  which  was  embraced  in  Avogadro's 
theory.  Inasmuch  as  the  density  of  oxygen  was  found  to 
be  about  sixteen  times  that  of  hydrogen,  it  will  be  further 


234  GENERAL  CHEMISTRY. 

noticed  that  the  theory  of  Berzelius  places  the  atomic 
weight  of  oxygen,  referred  to  hydrogen,  as  16  :  1  instead  of 
8  :  1,  as  in  the  Dalton  system,  the  former  assuming  two  atoms 
of  hydrogen  with  one  atom  of  oxygen  in  the  water  mole- 
cule while  the  latter  assumed  one  atom  of  each. 

To  find  the  atomic  weights  of  the  nongaseous  elements 
from  their  combining  proportions  Berzelius  announced  an- 
other principle,viz.:  that  in  all  compounds  of  two  elements  one 
must  be  present  as  a  single  atom.  This  rule,  it  will  be  recog- 
nized, is  perfectly  arbitrary,  resting  on  no  experimental 
basis.  Yet  by  combining  his  two  principles  he  was  able 
to  make  many  valuable  determinations,  some  of  which  are 
still  recognized  as  practically  correct.  In  the  cases  of  most 
of  the  metals,  however,  his  first  published  atomic  weights 
were  twice  as  great  as  those  now  commonly  accepted,  and 
in  fact  twice  as  great  as  the  weights  he  announced  later 
himself.  It  should  be  said  here  that  most  of  our  present 
atomic  weight  values  are  based  on  the  original  determina- 
tions of  Berzelius.  It  will  be  interesting  to  note  the  figures 
actually  given  by  this  writer  in  1815  and  1826.  In  calcu- 
lating his  first  table  he  made  certain  assumptions  regarding 
the  composition  of  the  oxides  which  led  him  to  the  double 
values  for  the  weights  of  the  metals.  In  ferrous  oxide,  for 
example,  he  assumed  a  composition  which  we  now  express 
as  FeO2.  Finding  in  this  for  200  parts  of  oxygen  693.6 
parts  of  iron,  he  considered  this  as  representing  the  weight 
of  one  atom  of  the  metal  on  his  oxygen  scale.  This  corre- 
sponds to  an  atomic  weight  of  111  on  our  present  system. 
But  giving  to  ferrous  oxide  the  composition  FeO,  the 
weight  would  be  reduced  one  half.  The  same  explanations 
apply  to  the  oxides  of  lead  and  copper  and  to  other  com- 
pounds which  Berzelius  worked  with.  Several  years  later, 
however,  a  new  principle  was  discovered  by  Mitscherlich, 
as  explained  below,  and  this  the  Swedish  chemist  accepted 
as  of  fundamental  importance.  He  applied  it  to  the  cor- 
rection of  his  weights,  giving  them  the  values  now  prac- 
tically retained  as  correct. 

In  the  following  table  some  of  the  Berzelius  atomic 
weights  of  1815  and  1826  are  given.  Those  of  the  latter 
year  are  reduced  to  the  hydrogen  scale,  as  shown  in  the 


GENERAL  CHEMISTRY. 


235 


fourth  column,  while  in  the  last  column  our  present  values 
are  given  for  comparison.  The  agreement  is  in  most 
cases  very  close. 


The  Weights   of  Berzelius. 


1815. 

1826. 

1826 

Reduced. 

Present 
values. 

Oxygen 

100 

100 

16  0 

16  0 

Sulphur 

201 

201  2 

32  2 

32  1 

Phosphorus  

167.5 

196.2 

31.4 

31  0 

Carbon 

74  9 

76  4 

12-5 

12  0 

Hydrogen                     

6  64 

6  24 

1  0 

1  0 

Arsenic 

839  9 

470  0 

75  2 

75  0 

Chromium  

708  1 

351  8 

56  6 

52  1 

Tellurium     ... 

806  5 

806  5 

129  1 

127  5 

Antimony. 

1618 

806  5 

129  1 

120  4 

Platinum  

1206  7 

1215  2 

194  5 

194  9 

Gold  

2483  8 

1243  0 

198  9 

197  2 

Mercury  .    ........ 

2531  6 

1265  8 

202  6 

200  0 

Silver  

2688.2 

1351  6 

216  3 

107  9 

Copper  .  . 

806.5 

395  7 

63  3 

63  6 

Nickel  

733  8 

369  7 

59.2 

58  7 

Lead  

2597  4 

1294  5 

207  1 

206  9 

Tin  

1470.6 

735  2 

117  6 

119  0 

Zinc  

806.4 

403  2 

64  5 

65  4 

343  0 

171  2 

27.4 

27  1 

Iron  

693  6 

339  2 

54  3 

56  0 

315.5 

158  4 

25.3 

24.3 

Calcium  

510.2 

256.0 

40.9 

40.1 

Strontium  

1118.1 

547.3 

87.6 

87.6 

Barium  

1709.1 

856.9 

137.1 

137  4 

Sodium 

579  3 

290  9 

46  5 

23  0 

Potassium  

978.0 

489.9 

78.4 

39.1 

It  will  be  observed  that  most  of  the  Berzelius  weights 
in  the  above  table  of  1826  agree  pretty  well  with  those  of 
the  present  time.  The  weights  of  sodium,  potassium  and 
silver  are  still  twice  as  great  as  the  modern  values,  how^ 
ever.  Berzelius  did  not  consider  nitrogen  and  chlorine  as 
simple  substances  and  hence  his  atomic  weights  for  them 
bear  no  simple  relation  to  any  now  recognized. 


236 


GENERAL  CHEMISTRY. 


Turner's  Table. 

Dr.  Edward  Turner,  a  contemporary  of  Daltonand  Ber- 
zelius,  published  not  a  little  on  the  subject  of  atomic 
weights.  His  results  were  partly  of  his  own  determination, 
but  more  largely  based  on  the  work  of  Dalton,  Wollaston 
and  Berzelius.  Wollaston,  like  Dalton,  took  the  relation 
of  hydrogen  to  oxygen  as  1:8,  but  assumed  O  =  10  as  the 
basis  of  his  system.  In  the  American  edition  of  Turner's 
Chemistry,  published  in  1822  and  1823,  the  author  gives  his 
own  table  in  detail,  with  H^l  as  the  basis  of  weights.  A 
part  of  this  table  is  here  quoted  for  comparison  with  that 
of  Berzelius. 

"TABLE  OF  CHEMICAL  EQUIVALENTS,  OR  ATOMIC  WEIGHTS." 

Hydrogen *f» 1 

Aluminum 18 

Antimony 44 

Arsenic 38 

Barium 70 

Bismuth 71 

Cadmium 56 

Calcium 20 

Carbon 6 

Chlorine 36 

Chromium 28 

Cobalt 30 

Copper 64 

Gold..                                      .  200 


Iron 28 

Lead 104 

Magnesium 12 

Mercury 200 

Nitrogen . 14 

Oxygen 8 

Phosphorus 12 

Potassium 40 

Silver 110 

Sodium 24 

Strontium 44 

Sulphur 16 

Tin 59 

Zinc 33 


It  will  be  noticed  that  many  of  the  numbers  in  the 
above  table  are  quite  different  from  those  given  by  Berze- 
lius, which  led  to  much  confusion  for  many  years.  It  is, 
however,  easily  recognized  that  by  multiplying  by  2  the 
weights  of  oxygen,  carbon,  sulphur,  chromium,  iron,  lead, 
magnesium,  strontium,  barium,  calcium,  tin  and  zinc  in 
the  Turner  table  values  very  close  to  those  of  the  Ber- 
zelius table  are  reached,  which  correspond  to  those  now  in 
use.  Also,  by  dividing  by  2  the  weights  for  sodium,  po- 
tassium and  silver  as  given  by  Berzelius,  we  obtain  num- 
bers agreeing  well  with  those  of  Turner,  accepted  at  the 
present  time.  It  is  evident,  therefore,  that  these  differ- 
ences may  be  traced  to  fundamental  differences  in  theories 


GENERAL  CHEMISTRY,  237 

of  combination  rather  than  to  differences  in  actual  results 
of  analyses.  With  the  proper  theory  of  the  union  of 
atoms  to  form  compounds,  it  was  recognized  by  these 
earlier  chemists  that  their  analytical  results  might  be  differ- 
ently interpreted  and  brought  into  accord.  Important 
aids  in  the  choice  of  atomic  weights  were  soon  forthcom- 
ing. 

THE  LAW  OF  ISOMORPHISM. 

In  1819  Mitscherlich  published  the  important  discovery 
that  in  many  groups  of  crystalline  compounds  of  like  form 
certain  elements  may  replace  each  other  without  causing 
essential  change  in  the  crystalline  structure.  Such  bodies 
he  described  as  isomorphous,  that  is,  of  like  form.  Thus, 
nickel,  cobalt,  zinc,  iron  and  magnesium  sulphates  are 
formed  by  the  solution  of  the  metals  or  oxides  in  sulphuric 
acid,  yielding  salts  which  resemble  each  other  in  form  and 
in  the  amount  of  water  of  crystallization  held.  In  another 
class  of  sulphates,  iron,  chromium  and  aluminum  replace 
each  other  perfectly.  These  metals  are,  therefore,  iso- 
morphous  with  each  other  in  their  compounds.  Many  of 
the  salts  of  phosphoric  and  arsenic  acids  are  found  to  crys- 
tallize in  similar  forms,  and  the  metals  in  them  may  be 
considered  as  isomorphous.  Now,  it  occurred  to  Berze- 
lius  and  others,  that  in  these  cases  the  metals  must  replace 
each  other  in  atomic  proportions,  and  given  the  atomic 
weight  of  a  metal  in  any  one  group  of  isomorphous 
substances,  the  other  atomic  weights  might  be  found 
by  a  determination  of  the  corresponding  replacing 
weights.  As  a  matter  of  fact,  Berzelius  employed  the 
principle. in  many  cases  tocorrect  his  earlier  determinations, 
and  it  is  still  of  value  in  choosing  between  possible  pro- 
portions found  by  analysis  alone. 

THE  LAW  OF  DULONQ  AND  PETIT. 

In  Chapter  II  a  unit  of  heat  was  defined  as  the  amount 
of  heat  required  to  raise  the  temperature  of  a  gram  of  water 
one  degree  centigrade.  It  was  known  to  physicists  before 


238  GENERAL  CHEMISTRY. 

the  end  of  the  last  century  that  very  different  amounts  of 
heat  are  required  to  warm  a  gram  of  other  substances 
through  the  same  range  of  temperature.  The  specific  heat 
of  a  substance  may  be  defined  as  the  fraction  of  a  unit  of 
heat  necessary  to  raise  the  temperature  of  one  gram  of  that 
substance  through  one  centigrade  degree.  These  specific 
heats  are  all  fractions,  with  one  exception,  that  of  hydro- 
gen. In  other  words,  less  than  a  unit  of  heat  is  sufficient 
to  raise  the  temperature  of  a  gram  of  the  substance  one 
degree.  For  many  of  the  metals  the  specific  heats  were 
known  with  a  fair  degree  of  accuracy  at  the  beginning  of 
this  century.  Working  further  on  the  determination  of 
these  data  two  French  physicists,  Dulong  and  Petit,  in 
1819,  made  the  interesting  discovery  that  the  observed 
specific  heats  seemed  to  vary  inversely  as  the  correspond- 
ing atomic  weights  as  determined  by  Berzelius.  For  many 
of  the  metals  the  product  of  the  atomic  weight  and  the 
specific  heat  was  found  to  be  practically  a  constant.  They 
found  further,  however,  that  to  bring  the  product  of  the 
atomic  weight  and  specific. heat  of  sulphur  as  given  by  Ber- 
zelius in  1815  to  agree  with  the  products  for  the  metals 
the  atomic  weights  of  many  of  the  latter  would  have  to  be 
reduced  to  one-half  the  figures  then  given.  Because  of 
this  suggestion  and  for  other  reasons  these  atomic  weights 
were  actually  reduced  and  so  appear  in  the  table 'of  1826. 
In  explanation  of  this  interesting  relation  of  atomic 
weights  to  specific  heats  Dulong  and  Petit  put  forth  the 
suggestion  that  the  atoms  have  all  the  same  capacity  for 
heat  and  that  therefore  the  specific  heat  of  a  mass  of  any 
given  substance  would  depend  on  the  number  of  atoms  in 
that  mass  and  consequently  on  the  atomic  weights;  of  the 
heavy  atoms  a  smaller  number  would  suffice  to  make  up  a 
given  weight  than  is  the  case  with  the  light  atoms.  To 
illustrate  this  the  following  short  table  is  appended,  in 
which  the  atomic  weights  and  specific  heats  as  now  every- 
where accepted  are  employed.  This  will  prove  more 
satisfactory  than  to  use  the  large  numbers  on  the  oxygen 
standard  of  Berzelius. 


GENERAL  CHEMISTRY. 


239 


ATOMIC 
WEIGHT. 

SPECIFIC 
HEAT. 

PRODUCT. 

120.4 

'        0    0523 

6  a 

Copper  .  . 

63  6 

0  .  0930 

5  9 

Iron  

56.0 

0  112 

6  3 

Magnesium  

24  3 

0  245 

6  0 

Manganese  

55  0 

0  122 

6  7 

Mercury  

200  0 

0  0319 

6  4 

Platinum  

194.9 

0  0325 

6  3 

Silver  

107.9 

0.0560 

6  0 

Thallium  

204.2 

0.0335 

6  8 

Zinc  

65.4 

0  09:<2 

6  1 

Lead  

206.9 

0  0315 

6.5 

Uranium 

239  6 

0  02bO 

6  7 

Gold  

197.2 

0  0324 

6  4 

Bromine  

80.0 

0  0843 

6  7 

Iodine.  . 

126.9 

0.0541 

6.8 

It  is  observed  that  the  product  is  nearly  a  constant, 
and  this  not  far  from  6.4.  With  the  large  numbers  of  Ber- 
zelius  the  product  would  be  about  40. 

If  it  maybe  assumed  that  this  relation  holds  good  for  all 
substances,  it  will  be  recognized  that  we  have  a  very  simple 
method  of  arriving  at  atomic  weights  which  cannot  be  fixed 
in  any  other  manner.  It  is  only  necessary  to  divide  the 
constant,  6.4,  by  the  experimentally  found  specific  heat  to 
reach  a  number  which  must  be  very  near  to  the  real  atomic 
weight.  For  instance,  the  specific  heat  of  thallium  is  .0335. 
This  divided  into  6.4  gives  191,  which  is  very  near  the 
atomic  weight  found  by  analysis.  More  will  be  said  about 
this  presently. 


THE  ELECTROLYTIC  EQUIVALENTS  OF  FARADAY. 

In  1834  Faraday  made  the  discovery  that  when  an  elec- 
tric current  passes  through  an  electrolyte  the  amount  of 
substance  separated  at  the  electrodes  depends  only  on  the 
intensity  of  the  current.  If  a  number  of  electrolytes  are 
placed  in  a  series  and  a  current  passed  through  the  whole, 
what  may  be  termed  equivalent  amounts  of  the  elements 
are  separated  in  the  same  time.  If  we  have  in  the  circuit 


240  GENERAL  CHEMISTRY. 

acidulated  water,  copper  sulphate,  silver  nitrate,  a  zinc 
and  an  iron  compound,  it  will  be  observed  that  in  the  time 
required  to  liberate  1  gram  of  hydrogen  from  the  water  the 
following  weights  of  the  metals  are  separated: 

Copper   31.6  grams. 

Silver 107.1  grams. 

Zinc 32.4  grams. 

Iron 27.8  grams. 

Such  results  were  obtained  for  a  large  number  of  ele- 
ments by  Faraday  and  others,  and  the  numerical  values 
found  were  by  many  looked  upon  as  representing  the  real 
atomic  weights.  But  difficulties  soon  appeared  when  it 
was  noticed  that  the  amount  of  separated  substance  de- 
pends in  some  cases  on  the  compound  from  which  it  is 
electrolyzed.  While,  for  instance,  31.6  Gin.  of  copper 
separates  from  the  sulphate,  63.2  would  be  liberated  from 
certain  other  compounds.  For  a  certain  class  of  iron  com- 
pounds the  iron  separated  amounts  to  27.8  Gin.,  but  for 
another  class  it  amounts  to  only  18.53  Gm.  in  the  same 
time.  It  is  not  clear  which  one  of  these  should  be  taken 
as  the  real  atomic  weight,  or  whether  either  should  be 
taken.  Other  instances  might  be  given  in  which  the  elec- 
trolytic equivalents  vary  with  the  nature  of  compounds 
from  which  obtained.  It  is  apparent,  therefore,  that  this 
method  of  fixing  atomic  weights  needs  itself  a  modification 
or  correction. 

By  this  time  it  is  evident  to  the  student  that  through 
the  multiplication  of  methods  proposed  for  fixing  the 
atomic  weights  great  confusion  must  have  resulted.  This 
was  indeed  the  case,  and  the  confusion  extended  to  the  for- 
mulas written  to  express  the  composition  of  many  simple 
compounds.  Water,  for  instance,  was  represented  by  four 
different  formulas  by  as  many  different  schools  of  chemists, 
and  the  reader  was  often  at  a  loss  to  understand  the  struc- 
ture of  a  body  as  expressed  by  symbols.  This  confusion 
existed  until  after  the  middle  of  the  century,  and  various 
suggestionsweremade  to  reconcile  conflicting  views  through 
a  more  rational  foundation  hypothesis. 


GENERAL  CHEMISTRY.  241 

THE  LAW  OF  AVOGADRO. 

It  was  stated  some  pages  back  that  Avogadro  was  among 
the  first  to  recognize  the  value  of  the  work  of  Gay  Lussac 
and  Humboldt  on  the  combination  of  gases  by  volume.  He 
was  the  first  to  draw  a  conclusion  of  the  highest  impor- 
tance from  them,  which  is  now  expressed  by  the  law  bear- 
ing his  name,  and  which  has  become  a  foundation  principle 
in  modern  scientific  chemistry.  The  law  of  Avogadro  may 
be  stated  in  this  way  :  Equal  volumes  of  all  gases  under  like 
conditions  of  temperature  and  pressure  contain  the  same  num- 
ber of  molecules.  The  statement  of  the  law  includes  the 
provision  that  the  gas  volumes  compared  must  be  taken 
at  identical  pressures  and  temperatures,  inasmuch  as 
changes  in  either  would  cause  corresponding  changes  in 
the  volumes,  and  therefore  in  the  number  of  molecules  in 
a  given  space. 

This  differs  from  the  conception  of  Berzelius  in  one 
important  particular.  The  latter  assumed  that  the  ultimate 
particles  of  the  simple  gases  exist  in  the  atomic  condition, 
and  that  equal  volumes  contain  the  same  number  of  these 
atoms.  But,  according  to  the  view  of  Avogadro,  these 
"  integrant  particles,"  or  molecules  may  be  and  usually  are 
made  up  by  the  union  of  several  atoms.  A  liter  of  hydro- 
gen, a  liter  of  oxygen,  a  liter  of  carbonic  acid  gas  and  a 
liter  of  ammonia  all  contain  the  same  number  of  molecules. 
It  must  be  said  here  that  the  proof  of  this  remarkable  law 
belongs  to  the  field  of  physics  rather  than  to  that  of  chem- 
istry. Its  truth  was  first  recognized  by  physicists,  espe- 
cially in  the  development  of  the  kinetic  theory  of  gases, 
and  was  later  accepted  by  chemists. 

It  follows  directly  from  the  law  of  Avogadro  that  molec- 
ular weights  of  simple  or  compound  gases  are  propor- 
tional to  their  densities.  In  the  following  table  the  densi- 
ties of  a  number  of  gases  referred  to  hydrogen  are  given 
and  in  the  last  column  the  molecular  weights  of  the  same 
gases  on  the  assumption  that  the  density  of  the  molecule 
of  hydrogen  referred  to  its  atom  is  2,  or  in  other  words 
that  the  molecule  of  hydrogen  contains  2  atoms.  The  rea- 
son for  this  assumption  will  follow. 


243 


GENERAL  CHEMISTRY. 


DENSITY. 

MOLECULAR 
WEIGHT. 

Hydrogen  

1.00 

2.0 

Oxvcen 

16  00 

32  0 

Water  vapor     

9  00 

18.0 

Chlorine         .    .    .        .        

35  50 

71  0 

Hydrochloric  acid   

18  25 

36  5 

80.00 

160  0 

40.50 

81  0 

Sulphur  (above   1000°) 

32  10 

64  2 

Hydrogen  sulphide                             

17  05 

34  1 

Nitrogen   

14.00 

28.0 

Ammonia                    

8  50 

17  0 

Carbon  monoxide                       

1400 

28.0 

Carbon  dioxide  

22.00 

440 

Marsh  gas  

8.00 

16.0 

The  following  considerations  indicate  why  it  is  believed 
that  the  molecules  of  hydrogen,  chlorine,  oxygen  and.  cer- 
tain other  elements  contain  two  atoms,  or  have  a  molecular 
weight  equal  to  twice  their  atomic  weights.  It  was  shown 
some  pages  back  in  what  proportions  certain  gases  com- 
bine by  volume.  Accepting  the  hypothesis  of  Avogadro 
as  correct,  we  may  substitute  for  volume  proportions  molec- 
ular proportions,  and  thus  reach  this  table  : 

1  molecule   of  hydrogen  -)-  i  molecule  of  chlorine  =  z  molecules  of  hydrochloric  acid . 

2  molecules  of  hydrogen  -f- 1  molecule  of  oxygen    =  z  molecules  of  water. 

3  molecules  of  hydrogen  -f-  i  molecule  of  nitrogen  =  z  molecules  of  ammonia. 

Now,  it  appears  that  from  the  one  molecule  of  chlorine 
we  obtain  chlorine  enough  to  form  two  molecules  of  hydro- 
chloric acid,  each  one  of  which  must  contain  both  chlorin*e 
and  hydrogen.  The  chlorine  molecule  must,  therefore, 
split  into  two  portions  to  do  this,  and  these  two  portions 
are  its  ultimate  atoms.  What  is  true  of  chlorine  is  also 
true  of  hydrogen.  The  oxygen  for  the  two  molecules  of 
water  vapor  comes  from  the  one  molecule  of  oxygen,  and 
the  nitrogen  for  the  two  molecules  of  ammonia  from  the 
one  molecule  of  nitrogen.  It  must  follow,  therefore,  that 
the  oxygen  and  nitrogen  molecules  are  also  double  because 


GENERAL  CHEMISTRY.  243 

they  divide  into  two  portions  in  going  into  combination. 
The  following  formulas  express,  then,  what  has  just  been 
said  : 


2H2+O2  =  2H2O 
3H2+N2  =  2H3N. 

It  would,  in  fact,  be  more  accurate  to  say  that  the  hy- 
drogen, oxygen,  chlorine  and  nitrogen  molecules  contain 
at  least  two  atoms.  Assuming  these  molecules  to  contain 
n  atoms  the  following  equation  would  also  be  true: 


We  have  here  a  splitting  into  two  parts  as  before.  But, 
granting  the  possibility  of  such  a  combination,  we  should 
expect  to  find  in  some  gaseous  hydrogen  compound  less 
than  half  as  much  hydrogen  as  we  find  in  what  we  call  the 
hydrogen  molecule.  However,  no  such  compound  has 
ever  been  discovered.  We  have  several  hydrogen  com- 
pounds which  contain  in  a  given  volume  just  half  as  much 
hydrogen  as  we  find  in  the  same  volume  of  pure  hydrogen 
gas,  but  no  compound  is  known  in  which  the  ratio  is  below 
this.  We  are  therefore  justified  in  placing  the  molecular 
weight  of  hydrogen  as  2  on  the  scale  which  has  the  atom 
of  hydrogen  as  its  basis,  and  from  this  the  other  molecular 
weights  in  the  table  above  follow. 

Suppose  now  that  we  apply  these  principles  to  the 
determination  of  the  atomic  weight  of  carbon.  What  we 
do  practically  is  to  find  the  smallest  weight  of  carbon  oc- 
curring in  the  molecular  weight  of  any  carbon  compound 
which  can  be  studied  in  the  gaseous  condition.  A  large 
number  of  such  compounds  are  known  and  analysis  shows 
that  the  smallest  amount  of  carbon  found  in  the  molec- 
ular weight  of  any  one  of  them  is  12.  Thus,  marsh  gas,  with 
a  molecular  weight  of  16,  contains  4  parts  of  hydrogen  and 
12  parts  of  carbon.  The  molecular  weight  of  carbon 
dioxide  is  44  and  we  find  here,  by  exact  analysis,  with  32 
parts  of  oxygen  12  parts  of  carbon.  The  molecular 
weight  of  carbon  monoxide  is  28,  and  this  we  find  is  made 
up  of  12  of  carbon  with  16  of  oxygen.  If,  however,  we 


244  GENERAL  CHEMISTRY. 

should  find  a  carbon  compound  with  only  6  parts  of  carbon 
in  the  molecular  weight,  as  derived  from  the  gas  density, 
we  should  be  obliged  to  conclude  that  12  represents  the 
weight  of  two  atoms.  In  reality,  therefore,  12  is  the  maxi- 
mum atomic  weight  which  may  be  assigned  to  carbon.  We 
are  obliged  to  admit  that  it  may  be  smaller,  although  the 
likelihood  of  this  is  very  remote  in  view  of  the  thorough 
study  already  given  to  carbon  compounds.  In  the  same 
manner,  from  their  combining  proportions  with  hydrogen 
and  oxygen  and  from  a  determination  of  molecular  weights 
of  their  compounds,  the  atomic  weights  of  the  other  ele- 
ments in  the  above  table  may  be  reached. 

Fortunately,  in  the  application  of  the  law  of  Avogadro 
to  the  fixing  of  atomic  and  molecular  weights  we  are  not 
limited  to  the  light  elements.  A  few  of  the  heavy  metals 
may  be  vaporized  under  conditions  which  permit  a  weigh- 
ing of  the  vapor,  and  for  a  still  greater  number  of  metals 
we  are  acquainted  with  compounds  which  are  volatile.  For 
zinc,  cadmium  and  mercury  the  vapor  densities  have  been 
determined  with  a  fair  degree  of  accuracy.  The  vapor  den- 
sities of  some  of  their  compounds  have  also  been  found, 
so  that  the  weight  of  the  atom  in  the  molecule  may  be 
determined.  Volatile  compounds  of  aluminum,  iron,  lead, 
thallium,  tin,  antimony,  bismuth,  chromium  and  several 
other  metals  have  also  been  studied  and  from  them  atomic 
weights  derived. 

The  credit  of  having  presented  the  wide  applications  of 
the  law  of  Avogadro  in  a  clear  and  simple  light  belongs 
largely  to  the  Italian  chemist,  Cannizzaro,  who  in  1858 
published  a  valuable  paper  on  the  subject.  In  this  he 
pointed  out  also  that  the  atomic  weights  derived  from  spe- 
cific heat  determinations  agree  with  those  found  from 
vapor  densities  in  a  number  of  cases  where  both  kinds  of 
experiments  could  be  made.  Since  his  time  many  new 
observations  have  been  made  which  confirm  -his  views 
fully. 

Having  now  two  methods  which,  when  applicable,  yield 
concordant  results,  it  is  possible  to  control  the  values 
reached  through  a  third  method,  the  observation  of  iso- 
morphism, for  a  large  number  of  cases.  Thus,  we  find  the 


GENERAL  CHEMISTRY.  245 

atomic  weight  of  zinc  from  its  vapor  density  and  specific 
heat.  As  isomorphous  with  zinc  we  have  the  metals  in  a 
well  characterized  series  of  sulphates.  The  amount  of 
metal  which  replaces  65.4  parts  of  zinc  in  these  sulphates 
we  look  upon  as  the  atomic  weights  of  these  metals.  For 
most  of  these  metals  we  are  not  able  to  determine  the 
atomic  weights  through  vapor  density,  but  specific  heat 
determinations  are  possible,  and  the  results  obtained  here 
are  in  perfect  accord  with  those  drawn  from  observations 
of  isomorphous  compounds. 

Three  general  methods  are  therefore  in  practical  use 
to  fix  atomic  weights.     These  are  : 

1.  From  the  molecular  weights  as  found  by  the  law  of 
Avogadro. 

2.  From  determination  of  specific  heats  of  bodies  in 
the  solid  condition. 

3.  From   a  determination  of  the  weights  of  the  sub- 
stances which   replace  each  other  in  isomorphous  com-, 
pounds. 

Some  illustrations  will  be  given  below. 


DETERMINATION  OF  VALENCE. 

In  the  sixth  chapter  the  general  notion  of  what  is  un- 
derstood by  valence  was  outlined,  but  nothing  was  said  as 
to  how  it  maybe  determined.  In  most  instances  a  knowl- 
edge of  the  atomic  weight  of  an  atom  is  necessary  to  fix  its 
valence.  In  certain  gaseous  compounds  it  follows  almost 
directly  from  the  manner  of  combination  of  their  elements. 

Fluorine,  chlorine,  bromine  and  iodine  form  only  one 
combination  with  hydrogen,  with  which  they  unite,  in  the 
gaseous  condition,  volume  for  volume  with  no  condensation. 
They  must  have,  therefore,  the  same  valence  as  hydrogen, 
which  arbitrarily  we  assume  to  be  unity,  as  we  haveno  knowl- 
edge of  the  real  nature  of  this  property  of  atoms.  Oxygen, 
sulphur  and  a  few  other  elements  can  combine  with  twice 
the  volume;  nitrogen,  arsenic  and  phosphorus  with  three 
times  the  volume;  and  carbon  and  silicon  with  four  times  the 


246  GENERAL  CHEMISTRY, 

volume  that  can  be  held  by  a  chlorine  atom.     The  follow- 
ing short  table  illustrates  these  types  of  combination: 

HF  H20  H3N  H4C 

HC1  H2S  H3P  H4Si 

HBr  H2Se  H3As 

HI  H2Te 

The  elements  in  the  first  column,  combined  with  hy- 
drogen, are  called  univalent;  those  in  the  second  column 
are  called  bivalent;  those  in  the  third  are  called  trivalent; 
those  in  the  fourth  quadrivalent. 

Elements  which  are  capable  of  replacing  hydrogen, 
atom  for  atom,  in  the  compounds  of  the  first  columns  are, 
like  hydrogen,  univalent,  and  this  is  determined  by  the 
observation  that  in  these  compounds  1  part  of  hydrogen  is 
replaced  by  weights  of  the  elements  represented  by  their 
atomic  weights.  Thus,  we  say  that  lithium,  sodium,  po- 
tassium and  silver  are  univalent  because  1  partof  hydrogen 
may  be  replaced  by  7,  23,  39.1  and  107.9  parts  respectively 
of  these  metals;  that  is,  by  amounts  corresponding  to  their 
atomic  weights.  The  elements,  calcium,  barium,  strontium 
and  magnesium  are  held  to  be  bivalent,  because  amounts 
corresponding  to  their  atomic  weights  replace  two  atoms  of 
hydrogen  from  two  molecules  of  the  compounds  in  the  first 
column,  or  two  atoms  of  hydrogen  in  one  molecule  of  the 
compounds  in  the  second  column  above.  Thus,  40.1  parts 
of  calcium,  or  1  atomic  weight,  combine  with  71  parts,  or 
2  atomic  weights  of  chlorine.  The  same  weight  of  cal- 
cium combines  with  16  parts,  or  1  atomic  weight  of  oxygen. 
In  the  same  way  aluminum  is  considered  a  trivalent  ele- 
ment because  27.1  parts,  representing  1  atomic  weight, 
combine  with  3  X  35.5  parts  of  chlorine.  Also,  2  X  27.1  parts 
of  aluminum  combine  with  3X16  parts  of  oxygen. 

In  general,  therefore,  the  valence  of  an  element  may  be 
fixed  by  a  study  of  the  compounds  it  forms  with  elements 
of  known  combining  capacity,  preferably  with  hydrogen, 
oxygen,  chlorine  or  nitrogen.  For  many  of  the  elements, 
it  has  been  discovered,  the  combining  capacity  is  not  con- 
stant, which  fact  is  illustrated  in  the  table  given  in 


GENERAL  CHEMISTRY.  247 

Chapter  VI.     As  yet  no  satisfactory  explanation  has  been 
offered  to  account  for  this. 

With  the  data  at  our  command  suppose  we  attempt  to 
determine  the  atomic  weight  of  some  element,  say  thallium. 
This  may  be  done  by  several  methods,  and  among  them 
by  rinding  the  amount  of  the  metal  in  combination  in  pure 
thallium  sulphate.  We  are  supposed  to  know  the  compo- 
sition of  various  sulphates.  Lead  sulphate  is  PbSO4, 
which  contains  206.9  parts  of  lead  to  96.1  parts  of  SO4, 
and  this  206.9  represents  the  atomic  weight  of  lead.  Potas- 
sium sulphate  is  K2SO4,  in  which  we  have  78.2  parts  of 
potassium  to  96.1  parts  of  SO4.  This  78.2  is  the  weight 
of  2  atoms  of  potassium,  as  we  discover  from  other  inves- 
tigations. Potassium  must  be  a  univalent  element,  while 
lead  is  bivalent.  Now,  when  we  examine  thallium  sulphate 
we  find  for  96.1  parts  of  SO4  408.4  parts  of  thallium.  If 
thallium  is  a  bivalent  element  like  lead  this  must  be  its 
atomic  weight.  If,  on  the  other  hand,  it  is  univalent,  like 
potassium  and  sodium,  408.4  must  represent  the  weight  of 
2  atoms,  giving  for  the  atomic  weight  204.2.  In  some 
respects  the  thallium  compounds  resemble  those  of  lead 
very  closely,  and  the  metals  themselves  are  much  alike.  In 
other  respects  thallium  should  be  considered  as  allied  to  the 
alkali  metals,  sodium  and  potassium.  It  dissolves  in  water 
at  a  high  temperature,  forming  an  alkaline  hydroxide.  The 
carbonate  is  also  soluble  and  alkaline,  while  that  of  lead  is 
insoluble.  We  have  here,  therefore,  an  illustration  of  a 
common  dilemma  which  presents  itself  in  the  practical 
determination  of  atomic  weights  from  combining  propor- 
tions, and  recourse  must  be  had  to  an  independent  method 
to  settle  the  question.  Notwithstanding  the  resemblance 
of  some  of  the  thallium  salts  to  those  of  lead,  the  analogies 
with  alkali  metal  compounds  are  closer  and  more  charac- 
teristic. Thus,  its  sulphate  is  soluble  and  isomorphous 
with  potassium  sulphate.  It  replaces  the  latter  in  common 
alum,  giving  a  thallium  alum,  isomorphous  with  potash 
alum.  It  forms  several  other  compounds  isomorphous 
with  potassium  compounds.  From  these  facts  it  would 
appear  that  thallium  should  be  taken  as  a  univalent  ele- 
ment, with  the  atomic  weight  204.2. 


248  GENERAL  CHEMISTRY. 

The  specific  heat  of  thallium  has  been  determined,  first 
by  one  of  the  discoverers  of  the  metal,  Lamy,  who  found  it 
to  be  0.0325.  Later,  a  more  accurate  determination  made 
it  0.0335.  From  the  law  of  Dulong  and  Petit  we  find 
the  atomic  weight  by  dividing  this  into  the  common  atomic 
heat,  6.4.  The  quotient  in  this  case  is  191,  which  is  a 
number  sufficiently  near  204.2  to  fix  this  as  the  atomic 
weight  rather  than  408.4. 

Few  of  the  thallium  compounds  are  volatile  enough  to 
admit  of  a  molecular  weight  determination  from  vapor 
density  according  to  the  law  of  Avogadro,  but  one  of  the 
chlorides  may  be  vaporized  without  decomposition.  Ros- 
coe  determined  the  vapor  density  of  thallous  chloride  and 
found  it,  referred  to  air.  to  be  8.2.  This  gives  for  the 
molecular  weight,  referred  to  hydrogen,  236.7,  which  is 
very  close  to  the  theoretical  molecular  weight  239.7,  cor- 
responding to  T1C1.  The  atomic  weight  of  thallium  must, 
therefore,  be  taken  as  204.2,  making  the  atom  univalent 
and  the  formula  of  the  sulphate  analyzed  T12SO4,  and  not 
T1SO4.  These  suggestions  illustrate  the  usual  procedure 
in  determinations  of  atomic  weights.  It  is  first  necessary 
to  find  by  exact  analysis  or  synthesis  the  combining  value 
of  the  body  in  question  with  some  element  or  group  of 
known  valence.  Then  it  remains  to  fix,  by  one  or  all  of 
the  three  methods  given,  this  combining  weight  as  the  real 
atomic  weight  or  as  a  multiple  of  it. 


CHAPTER  XII. 


CLASSIFICATION    OF    THE    ELEMENTS.— GENERAL 

PROPERTIES  OF  THE  METALS  AND 

THEIR  SALTS. 

IN  THE  preceding  chapters  some  of  the  most  important 
elements  with  a  few  of  their  compounds  have  been 
described.  The  order  in  which  the  several  elements  were 
considered  was  a  practical  rather  than  a  scientific  one, 
being  determined  largely  by  their  abundance  and  general 
usefulness  in  the  arts  in  some  cases,  and  in  others  by 
peculiarities  in  the  classes  of  compounds  formed  with  them. 
These  elements  are  frequently  spoken  of  as  the  negative  or 
nonmetallic  elements.  Most  of  them  combine  readily  with 
the  so  called  positive  or  metallic  elements,  which  remain  to 
be  described.  Before  taking  up  a  discussion  of  the  metals 
a  systematic  classification  of  the  elements,  as  a  whole, 
must  be  given  and  that  will  follow  here. 

From  the  very  nature  of  the  case  it  is  evident  that  any 
classification  attempted  must  be  in  many  respects  arbi- 
trary. The  common  grouping  of  the  elements  as  metals 
and  nonmetals  is  an  illustration  of  an  arbitrary  division 
which,  while  useful  in  some  directions,  is  very  unsatisfac- 
tory in  others,  for  the  simple  reason  that  there  is  no  sharp 
line  of  demarcation  between  the  two  groups.  Many  sub- 
stances can  be  considered  just  as  well  in  one  group  as  in 
the  other.  In  the  preceding  pages  the  student  has  seen 
that  certain  elements  were  naturally  thrown  together.  Thus 
fluorine,  chlorine,  bromine  and  iodine  form  a  natural  fam- 
ily and  have  been  so  considered  by  chemists  for  many 
years.  Sulphur,  selenium  and  tellurium  are  similarly 
treated  because  in  their  properties  and  in  the  compounds 


250  GENERAL  CHEMISTRY. 

they  form  they  show  marked  resemblances.  Among  the 
metals  similar  related  substances  will  be  found.  The  fol- 
lowing table  shows  several  families  or  groups  of  triads  to 
which  attention  was  early  called  by  different  chemists. 

The  names    of   the  elements  and    the   corresponding 
atomic  weights  are  given: 


Lithium  .  .  . 

7.03 

Calcium..  . 

40.07 

Sulphur  .  . 

....   32  07 

Sodium.  .  .  . 

.  ..     23.05 

Strontium 

..    ..  87.61 

Selenium  . 

....  79.02 

Potassium.. 

39.11 

Barium  . 

..137.43 

Tellurium 

..127.49 

Chlorine 

35 

45 

Iron 

.    .  56 

0? 

|  Osmium 

1!)0 

(»q 

Bromine 

79 

95 

Nickel    .  .  . 

58 

r.q 

|  Iridium 

193 

V>. 

Iodine  

..  126 

85 

Cobalt  

....  58. 

93 

|  Platinum  .  . 

...194 

89 

The  properties  of  the  elements  in  these  groups  vary 
with  changes  in  the  atomic  weights.  In  the  iron  group  the 
weights  are  nearly  the  same  and  we  find  that  the  metals 
and  their  compounds  are  much  alike.  The  same  is 
true  in  the  platinum  group.  A  consideration  of  such  rela- 
tions would  seem  to  suggest  that  the  properties  of  ele- 
ments may  depend  on  their  atomic  weights,  and  this  has 
been  shown  to  be  in  a  marked  degree  the  case. 

THE  PROPERTIES  OF  THE  ELEMENTS  AS  PERIODIC 
FUNCTIONS  OF  THEIR  ATOMIC  WEIGHTS. 

In  order  to  show  any  relations  existing  between  the 
atomic  weights  and  properties  of  elements  let  us  write  their 
symbols  in  the  order  of  increasing  weights,  beginning  with 
lithium,  as  follows: 

Li,  7.03;    Be,  9.08;    B,  10.95;   C,  12.01;   N,  14.04;    O,  16.00;    F,  19.08. 

We  have  here  a  gradual  increase  in  the  atomic  weights 
and  a  well  characterized  change  in  properties  correspond- 
ing. Lithium  is  strongly  metallic  and  positive  in  its 
behavior,  while  fluorine  is  as  certainly  nonmetallic  and 
negative.  The  next  greatest  atomic  weight  is  that  of 
sodium  =  23. 05.  But  we  have  here  an  element  with  prop- 
erties like  those  of  lithium  rather  than  like  those  of 
fluorine.  Sodium  evidently  does  not  follow  the  latter  ele- 


GENERAL  CHEMISTRY.  251 

ment  in  the  series  as  begun.     Let  us  therefore  make  a  new 

series,  parallel  with  the  first,  which  runs  thus: 

Na,  23.05;  Mg,  24.28;  Al,  27.11;  Si,  28.40;  P,  31.02;  S,  32.07;  Cl,  35.45. 

We  have  a  new  series  of  seven  elements,  beginning  with 
a  characteristic  metal  and  ending  with  a  characteristic 
nonmetal.  The  corresponding  elements  in  the  two  series 
resemble  each  other  very  closely  in  their  chemical  behav- 
ior and  in  the  compounds  they  form.  It  appears  from  this 
that  certain  properties  are  repeated  in  passing  through  a 
series  or  period  of  seven  elements,  involving  a  change  in 
atomic  weight  of  about  16  units.  What  has  been  done 
here  for  fourteen  elements  may  be  done  for  many  more. 
This  grouping  was  suggested  nearly  30  years  ago  by  Lothar 
Meyer  and  D.  Mendelejeff,  independently,  and  is  called  the 
Periodic  Arrangement  of  the  elements.  It  is  shown  in  the 
following  table  in  which  the  symbols  and  atomic  weights 
are  given,  and  also  the  differences  in  atomic  weights  in 
passing  from  one  period  to  another.  Some  of  the  elements 
cannot  be  well  included  in  the  seven  families  or  groups  as 
first  indicated  and  are  therefore  placed  in  a  separate  or 
eighth  group. 

This  periodic  arrangement  of  the  elements  is  often 
called  the  natural  arrangement,  but  the  appropriateness  of 
this  term  may  not  be  immediately  apparent.  The  close 
relations  of  the  elements  of  the  first  group,  Li,  Na,  K,  Rb 
and  Cs  and  their  compounds  are  easily  recognized,  but  the 
position  of  the  other  elements,  Cu,  Ag  and  Au  is  not  as 
evident.  In  studying  the  compounds  of  the  metals  later 
the  student  will  find  that  the  strongest  likenesses  are  found 
between  these  compounds  rather  than  between  the  metals 
themselves.  In  the  next  group  the  metals  Be,  Mg,  Ca,  Sr 
and  Ba  are  very  closely  related,  while  among  their  com- 
pounds there  are  the  closest  resemblances.  The  salts  of 
Zn  and  Cd  are  in  many  instances  much  like  those  of  Mg, 
and  hence  the  propriety  of  grouping  them  all  together. 
The  evidence  for  the  position  of  Hg  is  not  as  clear.  In  the 
third  and  fourth  groups  there  are  a  number  of  rare  ele- 
ments which  are  not  as  easily  compared;  but  in  the  fifth 
group  N,  P,  As,  Sb  and  Bi  form  a  line  as  they  should  on 


252 


GENERAL  CHEMISTRY. 


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GENERAL  CHEMISTRY.  253 

account  of  the  related  compounds  they  yield  with  several 
other  elements. 

In  the  sixth  group  we  have  O,  S,  Se  and  Te,  the  rela- 
tions of  which  have  been  already  pointed  out.  Also,  as  a 
subgroup,  Cr,  Mo,  W  and  U  are  naturally  and  closely  re- 
lated. In  the  seventh  group  the  analogies  of  F,  Cl,  Br 
and  I  are  well  marked  and  already  known  to  the  student. 
The  position  of  Mn  may  be  questioned,  as  we  think  of  this 
element  as  characteristically  metallic.  But,  like  chro- 
mium, we  have  here  an  element  which  often  acts  in  the 
nonmetallic  condition  forming  salts  known  as  manga- 
nates  and  permanganates,  which  latter  resemble  the  per- 
chlorates  in  some  respects,  and  this  in  a  measure  justi- 
fies its  position  as  a  part  of  the  group. 

It  will  be  noticed  that  a  number  of  vacant  spaces  are 
left  in  the  table.  As  first  published  spaces  were  left 
after  Ca  and  Zn.  Mendelejeff  predicted  the  discovery  of 
elements  to  fill  these  places  and  also,  approximately,  their 
atomic  weights  and  physical  properties.  Later  scandium, 
gallium  and  germanium  were  found  in  rare  minerals  and 
investigation  showed  that  they  would  occupy  the  vacant 
spaces  very  well.  Some  of  the  spaces  still  empty  may  yet 
be  filled  by  known  elements  after  a  more  thorough  study 
of  their  properties.  A  number  of  rare  metals  are  known 
to  which  atomic  weights  between  about  140  and  180  have 
been  given.  Further  investigation  is  needed  to  settle 
fully  the  true  weights  of  these  substances  and  their  prop- 
erties. 

Notwithstanding  the  evident  shortcomings  of  the  table, 
it  presents  as  a  whole  the  most  consistent  arrangement 
of  the  elements  yet  attempted  and  in  the  following  chap- 
ters they  will  be  studied  in  families  or  groups  as  shown. 
The  student  will  then  see  relations  more  clearly  than 
can  be  pointed  out  in  this  brief  introduction. 

ATOMIC  VOLUMES. 

Some  interesting  relationships  are  disclosed  when  we 
compare  the  relative  volumes  occupied  by  an  atom  of  each 
of  the  metals.  These  relations  may  be  best  shown  by  com- 


254 


GENERAL   CHEMISTRY. 


paring  the  volumes  filled  by  atomic  weights  of  the  different 
elements  expressed  in  grams.  The  term  atomic  volume 
is  applied  to  the  quotient  obtained  by  dividing  the  atomic 
weight  in  grams  by  the  specific  gravity  of  the  element. 
Thus,  for  iron  we  have  -£1  =  7.2.  That  is,  the  atomic 
weight  of  iron  in  grams  occupies  7.2  cubic  centimeters. 
Whether  we  use  grams,  milligrams  or  millionths  of  milli- 
grams as  our  standard  of  weight,  the  relations  between  the 
atomic  volumes  calculated  must  remain  the  same.  In  the 
following  chapters  the  numerical  values  of  the  atomic  vol- 
umes will  be  given  and  their  relations  shown. 


THE  VALENCE  OF  THE  GROUPS  IN  THE  PERIODIC  SYSTEM. 

It  has  been  pointed  out  that  the  maximum  valence  of 
most  of  the  elements  is  shown  by  the  number  of  the  group 
to  which  they  belong.  There  are  a  few  exceptions  to  this 
statement,  but  the  general  truth  of  it  is  so  striking  that  a 
table  in  illustration  of  the  facts  is  here  appended.  The 
valence  is  measured  by  the  oxygen  combining  power  of 
the  atoms,  and  for  the  purpose  of  better  comparison  the 
formulas  in  some  of  the  groups  are  doubled,  so  as  to  have 
always  two  atoms  of  each  element  in  combination  with  the 
oxygen.  Thus,  lime  is  written  Ca2O2,  instead  of  CaO, 
as  usual. 


I 

II 

III 

IV 

V 

VI 

VII 

VIII 

Li2O 

Be-jOs 

B2O, 

C2O4 

N2O5 

Na2O 
K2O 
Cu2O 

Mg202 
Ca202 
Zn2O2 

A1203 

Sc203 
Ga2O3 

Si2O4 
Ti204 
Ge2O4 

PS05 
V205 
As2O6 

S206 
Cr206 
Se2O6 

C1207 
Mn2O7 
Br2O7 

Rb2O 

Sr«O« 

Yt2O8 

Zr2O4 

Cb2O6 

Mo2O6 

Ru2O8 

Ag80 
POOO 

Cd2O2 
BaoOo 

In2O8 
La,Oo 

Sn2O4 
CeoCh 

Sb205 
Ta2Oe 

Te206 
W2O6 

I20, 

Os2Oa 

AiioO 

HeoOo 

TloO-, 

PboO* 

Bi2Os 

U2O6 

Among  the  exceptions  referred  to  these  should  be  men- 
tioned. Bivalent  compounds  of  copper  are  well  known, 
and  the  oxide,  CuO,  is  a  common  substance.  Some  of  the 


GENERAL  CHEMISTRY.  255 

most  stable  compounds  of  gold  are  trivalent,  instead  of 
univalent.  In  the  eighth  group  the  higher  oxygen  com- 
pounds are  not  well  characterized,  and  are  therefore 
omitted.  Otherwise,  the  table  shows  highly  instructive 
relations  in  condensed  form  and  should  be  thoroughly 
studied. 

THE  METALLIC  ELEMENTS. 

As  already  intimated,  the  distinction  between  metallic 
and  nonmetallic  elements  is  not  a  very  sharp  one,  but  is 
observed  in  a  general  way  as  a  matter  of  convenience.  The 
metallic  characteristics  of  all  the  elements  of  the  first  and 
second  groups  in  the  Periodic  system  are  well  defined  and 
easily  observed.  At  the  beginning  of  the  third  group  we 
find,  however,  an  element  which  behaves  in  general  as  a 
nonmetal.  The  other  elements  of  the  group  show  pro- 
nounced metallic  properties.  In  the  fourth  group  carbon 
and  silicon  are  pronounced  nonmetals,  while  the  others 
maybe  considered  in  general  as  metals.  The  first  elements 
in  the  fifth  and  sixth  are  nonmetals,  while  the  others  are 
more  distinctly  metallic.  Arsenic  illustrates  well  the  diffi- 
culty in  drawing  a  close  distinction  between  the  two  classes 
of  elements.  In  its  physical  appearance  it  is  metallic, 
resembling  cast  iron;  but  on  the  other  hand,  in  its  most 
characteristic  combinations  it  forms  the  nonmetallic  or  acid 
part.  In  the  seventh  group  all  the  elements,  with  the  ex- 
ception of  manganese,  are  pronounced  nonmetals.  In  the 
eighth  group  all  the  elements  are  properly  metals, 
i  From  a  merely  superficial  point  of  view  we  distinguish 
certain  properties  which  are  spoken  of  as  metallic  properties. 
The  most  characteristic  of  these  are  luster,  high  specific 
gravity,  hardness,  conducting  power  for  heat  and  electricity, 
malleability,  ductility,  and  capability  of  forming  alloys.  But 
few  of  the  metals  possess  all  of  these  properties  in  a 
marked  degree. 

Alloys  are  combinations  formed  between  metals,  in 
which,  as  a  rule,  the  combining  proportions  bear  no  rela- 
tion to  atomic  weights.  Thus,  brass  is  an  alloy  of  cop- 
per and  zinc  in  the  proportion,  2  :  1,  approximately,  while 
the  atomic  weights  are  63.6  and  65.4.  An  amalgam  is  an 


256  GENERAL  CHEMISTRY. 

alloy  in  which  one  of  the  metals  is  mercury.  Among  the 
common  alloys  there  may  be  mentioned  brass,  solder,  type 
metal,  pewter,  German  silver,  bronze  and  the  mixtures 
employed  in  coinage  and  jewelry. 

In  alloys,  as  a  rule,  the  melting  point  is  below  that 
which  would  be  calculated  from  the  proportions  and  melt- 
ing points  of  the  constituents  taken.  In  some  cases  it  is 
below  that  of  the  most  readily  melted  of  its  constituents. 
Fusible  metals  and  very  soft  solders  are  made  of  such  mix- 
tures, some  of  which  will  be  described  later. 


IMPORTANT  CLASSES  OF  METALLIC  COMPOUNDS. 

The  student  is  already  familiar  with  the  fact  that  com- 
pounds are  usually  formed  by  combination  of  elements 
of  opposite  chemical  nature.  In  Chapter  IV  it  was  shown 
that  salts  are  formed  by  the  union  of  metals  with  non- 
metals  or  by  the  action  of  metals  on  acids.  Salts  consti- 
tute a  very  important  class  of  inorganic  compounds.  The 
oxides  and  hydroxides  of  the  metals  are  also  important 
and  numerous. 

The  following  brief  definitions  may  be  found  useful  to 
the  student: 

Acids  are  compounds  formed  by  the  union  of  hydrogen 
with  nonmetallic  elements,  or  groups,  the  hydrogen  being 
replaceable  by  metals  to  form  salts.  Acid  oxides  or  anhy- 
drides are  combinations  of  oxygen  with  nonmetallic  ele- 
ments, which  unite  with  water  to  form  acids.  Thus,  SO3 
unites  with  water  to  form  H2SO3. 

Bases  are  combinations  of  metallic  elements  with 
hydrogen  and  oxygen  united  as  hydroxyl.  Thus,  KOH  and 
A1(OH)3  are  bases.  The  bases  which  are  soluble  in  water 
are  alkalies.  KOH,  NaOH  and  Ba(OH)2  are  illustrations. 
Basic  oxides  or  anhydrides  are  combinations  of  metallic 
elements  with  oxygen ;  they  yield  bases  by  combination  with 
water,  and  form  salts  by  combination  with  acids  or  acid 
anhydrides.  Thus,  CaO  is  a  basic  oxide;  with  water  it 
forms  Ca(OH)8,  a  true  base.  With  SO3  it  forms  CaSO4. 


GENERAL  CHEMISTRY.  257 

Salts  are  compounds  formed  by  the  union  of  metallic 
with  nonmetallic  elements.  Practically  they  may  be  made 
by  combination  of  the  metal  with  the  nonmetal  directly, 
as  Na-f-Cl  =  NaCl;  or  by  the  action  of  an  acid  on  the  metal, 
as,  Zn-f-H3SO4  — ZnSO4-f-H2;  or  by  the  union  of 
an  acid  with  a  base,  as,  HCl-fKOHr=KCl-|-H2O;  or  by 
the  union  of  a  basic  oxide  with  an  acid  oxide,  as 
K2O-|-SO3  =  K2SO4.  Salts  are  formed  also  by  a  variety 
of  indirect  methods,  but  these  illustrations  sufficiently 
characterize  their  composition. 

The  following  list  embraces  the  most  important  classes 
of  metallic  compounds,  some  of  which  will  be  but  briefly 
considered  in  discussing  the  metals: 

1.  Oxides  and  hydroxides. 

2.  Sulphides. 

3.  Carbonates. 

4.  Halogen  salts,   or   binary  compounds   of  fluorine, 

chlorine,  bromine  and  iodine. 

5.  Compounds    with  the  acids  of  sulphur;   sulphites 

and  sulphates. 

6.  Compounds  with  the  acids  of  nitrogen;  nitrites  and 

nitrates. 

7.  Compounds  with  the  oxygen  acids  of  the  halogens; 

hypochlorites,  chlorates,  etc. 

8.  Borates,   phosphates,    arsenites,    arsenates,    chro- 

mates,  manganates  and  silicates. 

Generalities  on  these  Compounds. 

Oxides.  The  important  groups  of  oxides  are  illus- 
trated in  the  short  table  above  in  which  variations  in  va- 
lence are  shown.  The  oxides  of  lithium,  sodium,  potassium, 
rubidium  and  caesium  dissolve  readily  in  water  forming 
alkaline  hydroxides.  The  oxides  of  calcium,  strontium 
and  barium  dissolve  slowly,  forming  alkaline  hydroxides. 
One  of  the  oxides  of  thallium,  T12O,  dissolves  easily,  form- 
ing a  hydroxide  like  those  of  sodium  and  potassium.  The 
highest  oxides  of  chromium  and  manganese  dissolve  in 
water,  forming  acids.  Practically  all  of  the  other  oxides  of 
the  metals  are  insoluble  in  water. 


258  GENERAL  CHEMISTRY. 

Hydroxides.  The  method  of  forming  some  of  the  hy- 
droxides has  just  been  explained.  Those  insoluble  in  water 
are  mostly  formed  by  precipitating  a  metallic  salt  by  a  sol- 
uble hydroxide,  thus: 


=  K2SO4-fZnO2H2. 

Some  of  these  hydroxides  redissolve  in  an  excess  of  the 
precipitant.  In  presence  of  strong  alkalies  they  appear  to 
behave  as  acids.  Thus  zinc  hydroxide  forms  potassium 
zincate: 

ZnO2H2+2KOH  =  K2O2Zn-f2H2O. 

Sulphides.  In  many  respects  the  sulphides  resemble 
the  oxides.  Only  those  of  the  light  metals  are  soluble  in 
water.  Many  of  them  may  be  formed  by  the  direct  union 
of  the  metal  with  sulphur  by  the  aid  of  heat.  The  sulphides 
of  the  heavy  metals  are  mostly  made  by  precipitation  of 
their  solutions  by  hydrogen  sulphide  or  ammonium  sul- 
phide. Thus, 


CuSO4  +  H2S  =  CuS+H2SO4 


The  behavior  of  the  sulphides  of  most  of  the  metals 
toward  acids  is  characteristic  and  important  in  analytical 
chemistry,  as  illustrated  in  a  preceding  chapter.  Many  of 
the  sulphides  dissolve  in  acids,  liberating  hydrogen  sul- 
phide just  as  the  corresponding  oxides  liberate  water: 


=  ZnCl2+H8S 
ZnO-f2HCl  = 


Several  of  the  heavier  sulphides  dissolve  in  solutions  of 
the  alkali^  sulphides,  forming  a  peculiar  class  of  salts  in 
which  sulphur  takes  the  place  of  oxygen.  Thus, 


which  is  called  potassium  thiostannate.  This  behavior  is 
important  in  analytical  chemistry,  and  is  employed  in  the 
separation  of  metals. 


GENERAL  CHEMISTRY.  259 

Carbonates.  The  carbonates  of  the  alkali  metals  are 
soluble  in  water.  The  others,  with  few  exceptions,  are 
insoluble.  All  carbonates  are  more  or  less  readily  decom- 
posed by  mineral  acids.  Two  general  classes  of  carbonates 
are  distinguished,  normal  and  acid,  or  bicarbonates.  As 
illustrating  the  first  kind,  we  have  common  sodium  car- 
bonate, Na2CO3.  As  illustrating  the  second,  we  have  the 
acid  carbonate,  commonly  called  bicarbonate,  HNaCO3. 
Bicarbonates  are  usually  made  by  the  action  of  carbonic 
acid  on  carbonates.  At  a  sufficiently  high  temperature 
carbonates  decompose,  with  formation  of  carbon  dioxide 
and  a  metallic  oxide.  As  further  illustrating  the  parallelism 
between  oxygen  and  sulphur,  it  may  be  mentioned  that  we 
have  a  class  of  bodies  called  sulphocarbonates,  or  thio- 
carbonates,  which  may  be  formed  by  reactions  analogous 
to  those  which  yield  carbonates: 

K2S+CS2  =  K2CS3 
K20-f-C02  =  K2C03. 

Acids  form  with  the  first  H2CS3,  which  breaks  down 
into  H2S  and  CS2.  The  second  yields  H2O  and  CO2 
almost  directly. 

Halides.  This  term  may  be  applied  to  the  binary 
compounds  of  the  metals  with  fluorine,  chlorine,  bromine 
and  iodine.  Of  the  formation  and  properties  of  the  fluo- 
rides little  need  be  said,  as,  with  few  exceptions,  they  are 
unimportant  substances.  The  halogen  compounds  of  most 
of  the  metals  are  soluble  in  water.  Of  the  insoluble  com- 
pounds those  of  silver,  lead  and  mercury  are  the  most 
characteristic.  The  halides  may  be  formed  by  a  number 
of  general  reactions,  of  which  the  following  are  the  most 
important: 

1.  By  direct   union   of  metals  with  the  halogen  ele- 
ment, as: 

Fe+Br2:=FeBr2. 

2.  By  action  of  a  metal  on  a  halogen  acid,  as: 


260  GENERAL  CHEMISTRY. 

3.    By  the  action  of  an  oxide  or  hydroxide  on  a  halogen 
acid,  as: 


KOH+  HC1  =  KC1   +  H3O. 

The  halides  of  most  of  the  metals  are  readily  decom- 
posed by  strong  sulphuric  acid,  with  formation  of  a  sul- 
phate. From  a  fluoride,  hydrofluoric  acid  is  produced,  and 
from  a  chloride,  hydrochloric  acid;  but  from  bromides  and 
iodides  some  free  bromine  and  iodine  are  obtained. 

Many  of  the  halides  are  volatile  at  a  high  temperature 
without  decomposition.  From  this  behavior  it  has  been 
found  possible  to  determine  their  vapor  densities  and 
therefore  molecular  weights.  In  the  halides  the  non- 
metallic  element  always  acts  with  unit  valence. 

Sulphates,  Sulphites  and  Allied  Salts.  It  has  been 
already  pointed  out  that  sulphur  forms  a  number  of  oxygen 
acids.  But  the  compounds  of  some  of  these  are  of  little 
importance  and  need  not  be  considered  here.  The  sul- 
phites, sulphates  and  thiosulphates  will  be  briefly  de- 
scribed. 

As  sulphurous  acid,  H2SO3,  is  a  compound  with  both 
hydrogens  replaceable  we  have  two  classes  of  salts  from  it, 
known  as  normal  and  acid  sulphites  or  bisulphites.  These 
have  the  general  formulas  M2SO3  and  MHSO3.  They 
are  characterized  by  their  instability  in  presence  of  oxidiz- 
ing agents  from  which  they  take  oxygen  and  become 
sulphates.  On  account  of  this  behavior  both  classes  of 
salts  are  used  in  photography  as  constituents  of  developers. 
The  sulphites  and  acid  sulphites  are  readily  decomposed 
by  acids  with  separation  of  SO2,  or  formation  of  H2SO3  in 
weak  solution.  The  normal  and  acid  sulphites  of  the 
alkali  metals  are  soluble  in  water;  the  normal  sulphites  of 
most  of  the  other  metals  are  either  insoluble  or  but  slightly 
soluble.  The  acid  sulphites,  however,  are  mostly  soluble. 
Their  solutions,  like  those  of  the  acid  carbonates,  decom- 
pose on  heating  with  separation  of  normal  sulphite  and 
escape  of  SO2. 

Sulphates  are  among  the  most  important  and  abundant 


GENERAL  CHEMISTRY.  261 

of  inorganic  salts.  They  are  most  readily  formed  by  solu- 
tion of  metals,  oxides  or  carbonates  in  dilute  sulphuric 
acid.  Acid  sulphates  of  the  type  MHSO4  are  known,  but 
with  the  exception  of  the  sodium  and  potassium  com- 
pounds they  have  little  importance.  Most  of  the  normal 
sulphates  are  soluble  in  water  and  dilute  acids.  The  sul- 
phates of  strontium  and  barium  are  practically  insoluble  in 
water  and  dilute  acids.  Calcium  sulphate  is  but  slightly 
soluble  in  water.  Lead  sulphate  is  likewise  but  little  sol- 
uble in  water.  This  behavior  is  applied  in  analytical  chem- 
istry in  the  separation  of  sulphuric  acid  and  also  of  the 
metals  lead,  strontium  and  barium. 

At   a  certain    temperature    acid    sulphates  of  the  type 
MHSO4  decompose,  yielding  pyrosulphates  and  water: 


At  a  higher  temperature  the  pyrosulphate  breaks  up 
into  normal  sulphate  and  SO3: 

M2S207  =  M2S04+S03. 

Attention  must  be  called  to  the  nomenclature  of  these 
salts.  Compounds  of  the  type  M2SO4  should  be  called 
sulphates  or  normal  sulphates;  those  of  the  type  MHSO4 
should  be  called  acid  sulphates,  but  are  sometimes  called 
bisulphates,  which  leads  to  confusion  with  the  next  group; 
compounds  of  the  type  M2S2O7  should  be  called  disul- 
phates  or  pyrosulphates. 

Of  the  pyrosulphates  only  those  of  sodium  and  potas- 
sium have  technical  importance. 

A  few  thiosulphates  are  practically  important.  These 
are  compounds  of  the  type  M2S2O3  and  may  be  looked 
upon  as  sulphates  in  which  one  atom  of  oxygen  is  replaced 
by  S. 

The  thiosulphates  of  the  alkali  metals  are  the  best 
known.  The  sodium  compound  is  made  in  large  quanti- 
ties as  described  in  a  former  chapter,  and  is  used  in  pho- 
tography. The  thiosulphates  are  decomposed  by  dilute 
acids  with  precipitation  of  sulphur,  and  liberation  of  sul- 
phurous oxide. 


262  GENERAL  CHEMISTRY, 

Nitrites.  These  salts  are  not  numerous;  only  those  of 
the  alkali  metals  have  importance.  Nitrites  are  made  by 
the  action  of  nitrous  acid  on  hydroxides, 

HN02+KOH  =  KN02+H20, 

or  by  the  reduction  of  nitrates.  When  an  alkali  nitrate  is 
heated  to  a  high  temperature  it  loses  part  of  its  oxygen. 
If  heated  with  lead  or  copper  it  loses  it  very  readily. 

KN03  =  KN02+0. 
Most  of  the  nitrites  are  very  soluble  in  water. 

Nitrates.  These  are  very  common  and  important  salts 
and  are  best  made  by  the  action  of  nitric  acid  on  metals, 
their  oxides,  hydroxides  or  carbonates. 

All  true  nitrates  are  soluble  in  water.  A  few  bodies 
known  as  basic  nitrates  are  insoluble.  All  nitrates  are  de- 
composed when  heated  with  sulphuric  acid,  and  all  suffer 
decomposition  by  heat  alone.  Some  yield  nitrites,  as  ex- 
plained above.  Some  break  down,  yielding  a  metallic  ox- 
ide, oxygen  and  nitric  oxide,  as  shown  by  the  behavior  of 
lead  nitrate  referred  to  in  chapter  VI.  All  nitrates  act  as 
strong  oxidizing  agents  when  fused  or  strongly  heated  with 
many  organic  substances. 

The  Oxygen  Salts  of  the  Halogens.  The  nomen- 
clature and  general  formulas  of  these  bodies  have  been 
illustrated  already.  The  acids  may  be  represented  by  the 
formulas  in  the  first  column  and  the  salts  by  the  formulas 
of  the  second, 

HOX         MOX 

HOXO        MOXO 

HOXO2       MOXO2 

HOX03       MOX03 

Salts  of  the  four  chlorine  oxygen  acids  are  known.  Hy- 
pobromites  and  bromates  are  also  known,  while  for  iodine 
the  iodates  and  periodates  are  the  only  stable  compounds. 
The  chlorine  salts  are  all  soluble  in  water ;  the  hypobro- 
mites  are  readily  soluble,  while  the  bromates  are  much 


GENERAL  CHEMISTRY.  263 

less  soluble  ;  the  iodates  of  the  alkali  metals  are  sparingly 
soluble  in  water.  Most  of  the  other  iodates  are  insol- 
uble. 

The  salts  of  all  of  these  acids  are  decomposed  by  heat, 
a  simple  halide  usually  resulting.  In  some  cases  the  de- 
composition takes  place  in  two  stages,  as  illustrated  by  the 
familiar  reaction  with  potassium  chlorate: 

2KC1O3  =  KClO4-fKCl-fO2 
KC104  =  KC1      +202. 

Sulphuric  acid  decomposes  all  the  salts  with  liberation 
of  acids  and  oxides.  Under  various  conditions  the  salts  are 
all  oxidizing  agents  ;  some  of  them  even  give  up  their  oxy- 
gen in  simple  aqueous  solution  to  bodies  easily  oxidized. 
This  is  well  illustrated  by  the  action  of  sodium  hypobro- 
mite  on  urea: 

3NaOBr-fCON2H4=3NaBr+C02-f2H2O+N2. 

Several  of  the  chlorates  and  hypochlorites  are  tech- 
nically important;  the  alkali  hypobromites  are  sometimes 
used  as  reagents,  as  illustrated  by  the  above  equation, 
while  the  other  salts  have  little  practical  value. 

Borates.  As  explained  in  the  eighth  chapter,  boric 
acid  forms  a  number  of  complex  borates,  which  is  possible 
from  the  peculiar  composition  of  the  acid.  Three  classes 
of  borates  are  known:  the  orthoborates,  M3BO3,  which 
are  not  very  stable,  the  metaborates,  MBO2,  of  which  a 
number  are  well  known,  and  the  pyroborates,  M2B4O7,  of 
which  common  borax  is  the  best  illustration.  The  alkali 
borates  are  soluble  in  water,  the  others  are  mostly  insolu- 
ble. All  borates  are  decomposed  by  strong  sulphuric  acid, 
with  liberation  of  boric  acid.  The  alkali  borates  may  be 
made  by  direct  union  of  boric  acid  with  carbonates  or 
hydroxides;  the  other  borates  are  usually  produced  by 
precipitation. 

Phosphates.  Three  principal  groups  of  phosphates, 
called  metaphosphates,  orthophosphates  and  pyrophos- 
phates  are  known.  The  first  are  of  the  type  MPO3,  the 


264  GENERAL  CHEMISTRY. 

molecule  of  metaphosphoric  acid  resembling  that  of  nitric 
acid.  Metaphosphates  seem  to  exist  in  condensed  or  poly- 
meric types  in  which  several  molecules  are  united,  as, 
M2P2O6,  M3P3O9,  M6PPO18.  The  salts  of  the  alkali  metals  only 
are  well  known;  some  are  soluble  in  water,  others  not;  all 
may  be  made  from  primary  orthophosphates  by  heat- 
ing and  cooling  carefully  under  varying  conditions.  By 
boiling  solutions  of  metaphosphates,  orthophosphates  are 
formed. 

The  orthophosphates,  or  common  phosphates,  are  rep- 
resented by  the  formulas  M3PO4,  M2HPO4  and  MH2PO4,  as 
phosphoric  acid  is  tribasic.  The  nomenclature  of  these 
salts  may  be  illustrated  by  the  sodium  compounds: 

Na3PO4  is  called  normal  sodium  phosphate,  trisodium 
phosphate  or  tertiary  sodium  phosphate. 

Na2HPO4  is  called  disodium  hydrogen  phosphate  or 
secondary  sodium  phosphate. 

NaH2PO4  is  called  dihydrogen  sodium  phosphate  or 
primary  sodium  phosphate. 

The  primary  phosphates  are  soluble  in  water  ;  the  sec- 
ondary and  tertiary  phosphates  are  mostly  insoluble,  those 
of  the  alkali  metals  being  exceptions.  The  tertiary  phos- 
phates are  very  stable  and  withstand  a  high  temperature. 
The  secondary  phosphates  are  converted  into  pyrophos- 
phates  by  heat  and  the  primary  phosphates  into  meta- 
phosphates. The  soluble  tertiary  phosphates  are  strongly 
alkaline,  the  secondary  phosphates  feebly  alkaline,  while 
the  primary  phosphates  are  acid  with  most  indicators. 

Pyrophosphates,  as  indicated  by  the  name,  are  com- 
monly made  by  heating  orthophosphates.  The  pyrophos- 
phates  exist  in  two  general  classes,  represented  by  the  for- 
mulas M4P2O7  and  M2H2P2O7.  The  secondary  ortho- 
phosphates  strongly  heated  give  the  first,  as 

2HNa2PO4— H2OrnNa4P2O7. 

Primary  phosphates  moderately  heated  give  the  second 
class,  as 

2H2NaPO4— H2O  =  Na2H2P2O7. 

The    alkali    pyrophosphates  are  soluble  in  water,  the 


GENERAL  CHEMISTRY.  265 

others  not.    Solutions  of  pyrophosphates  when  heated  with 
acids  yield  orthophosphates. 

A  few  phosphites  and  hypophosphites  are  known,  but 
as  they  are  comparatively  unimportant,  they  need  not  be 
discussed  here. 

Arsenites.  These  salts  have  the  general  formula 
M3AsO3.  Those  of  the  alkali  metals  may  be  prepared  by 
dissolving  arsenous  oxide  in  solutions  of  hydroxides  or  car- 
bonates. These  products  are  soluble  in  water.  The  arsen- 
ites  of  the  heavy  metals  are  insoluble  and  are  usually 
made  by  precipitation  of  the  sodium  or  potassium  salt  by 
a  salt  of  a  heavy  metal.  Several  metarsenites  are  known 
with  the  general  formula  MAsO2. 

Arsenates.  Compounds  of  the  general  formula 
M3AsO4  are  made  by  dissolving  arsenic  oxide  in  alkali 
solutions.  Crude  products  are  formed  also  by  fusing  a  mix- 
ture of  arsenous  oxide,  niter  and  alkali.  The  arsenates 
are  isomorphous  with  the  phosphates  and  resemble  them 
in  many  respects.  The  normal  arsenates  of  the  alkali 
metals  are  soluble  in  water,  the  other  arsenates  are  insol- 
uble. Acid  arsenates  corresponding  to  acid  phosphates 
are  known.  These  have  the  general  formulas  MH2AsO4 
and  M2HAsO4.  Hydrogen  sulphide  reduces  acidified  so- 
lutions of  arsenates  slowly,  separating  sulphur  and  pre- 
cipitating As2S3. 

Chromates.  The  salts  of  chromic  acid  have  the  for- 
mula M2CrO4,  corresponding  to  the  sulphates.  The  chro- 
mates  of  the  alkali  metals  are  soluble  and  are  usually  made 
directly  from  the  common  ore,  known  as  chrome  iron- 
stone, by  fusion  with  alkali  carbonate  in  presence  of  air. 
The  chromates  of  the  other  metals  are  insoluble  in  water 
and  are  made  by  precipitation,  as, 

4  +  PbN206=PbCr04 


By  the  action  of  acids  chromates  are  converted  into 
dichromates  or  anhydrochromates,  of  the  type  M2Cr2O7, 
thus: 


266  GENERAL  CHEMISTRY. 

The  dichromates  are  usually  red,  while  the  chromates 
are  mostly  yellow.  The  dichromates  are  active  oxidizing 
agents,  those  of  sodium  arid  potassium  being  valuable 
mainly  on  this  account.  In  presence  of  sulphuric  acid  and 
reducing  bodies  they  give  up  oxygen  in  this  way: 

K2Cr,O7  +  4H2SO4:=K2SO4+Cr2(SO4)3+4H2O4-3O. 

It  will  be  recalled  that  the  dichromate  was  employed  in 
the  liberation  of  chlorine  from  hydrochloric  acid.  This  is 
also  an  oxidation  reaction  : 

K8Crg07  +  14HCl  =  2KCl+2CrCls+7H80+3Cl8. 

In  both  cases  the  acid  combinations  are  destroyed  and 
salts  of  chromium  result. 

Manganates.  By  fusion  of  powdered  black  oxide  of 
manganese  with  caustic  alkali  in  presence  of  oxidizing 
agents,  or  the  air  even,  a  green  manganate  results, 
M2MnO4.  Such  salts  are  not  stable  and  are  not  important. 
By  action  of  acids  the  manganates  become  converted  into 
the  permanganates,  of  which  the  potassium  compound  is 
the  best  known.  It  has  the  formula  KMnO4.  These  salts 
are  important  because  of  the  fact  that  they  give  up  oxygen 
readily,  and  are  extremely  active  oxidizing  agents.  Like 
the  dichromates,  they  become  reduced  to  salts  of  the  active 
element  in  the  decomposition: 

2KMnO4  +  3H2SO4  =  K2SO4+2MnSO4+3H2O+5O. 

The  permanganates  are  all  highly  colored  compounds, 
the  common  soluble  salts  being  deep  purple. 

Silicates.  In  Chapter  VIII  some  of  the  important 
properties  of  silicic  acid  were  explained  and  formulas  of 
silicates  given.  Very  few  of  the  silicates  can  be  made  in 
pure  condition  and  our  knowledge  is  chiefly  confined  to 
mixtures  obtained  by  fusion  of  silica  with  metallic  oxides 
or  carbonates,  or  to  crude  precipitation  products  formed 
by  adding  salt  solutions  to  solutions  of  water  glass.  The 
silicates  and  borates  are  alike  in  this  respect,  that  they 


GENERAL  CHEMISTRY.  267 

exist  in  very  complex  condensed  types  which  may  be 
looked  upon  as  derived  from  acids  formed  by  the  loss  of 
one  or  more  molecules  of  water  from  several  molecules  of 
the  ortho  acid.  Common  borax  is  the  sodium  salt  of  the 
acid  related  to  the  ortho  acid  in  this  manner: 

4H3B03— 5H20  =  H2B407. 

Many  silicates  are  similarly  derived. 
The  salts  of  orthosilicic  acid  are  called  orthosilicates. 
Those  from  the  first  derived  acid  are  called  metasilicates. 
Salts  derived  from  condensed  silicic  acids  are  called  di-, 
tri-,  or  polysilicates  in  general.  The  following  formulas  illus- 
trate this: 

Orthosilicates  M4SiO4. 

Metasilicates  M2SiO, 


Disilicates  M6Si2O7. 

M2Si206. 
Trisilicates  M4Si3O8. 

M10Si3011. 

SPECIFIC  HEATS  OF  COMPOUNDS. 

For  a  number  of  compounds  it  has  been  shown  that  the 
molecular  heat  is  practically  the  sum  of  the  atomic  heats. 
To  illustrate  this  it  will  be  recalled  that  the  .following 
atomic  heats  are  known  and  given  in  the  last  chapter: 
Lead,  6.5;  silver,  6.0;  bromine,  6.7;  iodine,  6.8.  From 
these  numbers  the  molecular  heats  could  be  calculated  as 
follows: 

PbI2,      6.5  +  2X6.8  =  20.1.  3X6.4  =  19.2 

PbBr2,  6.5-1-2X6.7  =  19.9.  3X6.4  =  19.2 

Agl,        6.0+        6.8  =  12.8.  2X6.4  =  12.8 

AgBr,    6.0+       6.7  =  12.7.  2X6.4  =  12.8 

By  actual  experiment  the  specific  heat  of  lead  iodide 
has  been  found  to  be  0.0427,  and  that  of  lead  bromide 
0.0533.  These  numbers,  multiplied  by  the  corresponding 
molecular  weights,  give  19.7  and  19.6  respectively  as  the 
molecular  heats.  For  silver  bromide  the  specific  heat  of 
0.074  has  been  found.  This,  multiplied  by  the  molecular 


268  GENERAL  CHEMISTRY. 

weight,  gives  13.9,  which  is  but  little  larger  than  the  cal- 
culated numbers  above.  For  silver  iodide  the  observed 
value  of  the  specific  heat  is  0  055,  which  gives  a  molecular 
heat  of  12. 9,  very  close  to  the  calculated  value.  From  such 
observations  it  has  been  found  possible  to  make  predic- 
tions concerning  the  numerical  values  of  specific  heats  in 
advance  of  their  actual  determination. 

Determination  of  Specific  Heat.    While  this  is  a 


FIG.  36. 

subject  which  properly  belongs  to  physics,  a  brief  explana- 
tion of  the  simple  principle  involved  in  practical  methods 
may  not  be  out  of  place.  Knowing  that  the  latent  heat  of 
the  melting  of  ice  is  79.5  units  we  can  find  the  specific  heat 
of  many  substances  which  are  insoluble  in  water  and  do 
not  decompose  water  by  observing  how  many  grams  of  ice 
are  melted  by  the  cooling  down  of  a  given  number  of 
grams  of  the  substance  in  question  through  a  given  range 
of  temperature.  To  make  this  test  the  metal  or  other 
body  under  experimentation  is  warmed  to  a  known  temper- 


GENERAL  CHEMISTRY.  269 

ature,  say  100°  C,  and  then  is  brought  immediately  in  con- 
tact with  an  excess  of  pure  ice  under  such  conditions  that 
all  its  heat  must  be  given  out  to  melt  part  of  it.  This  may 
be  done  by  the  aid  of  an  apparatus  called  an  ice  calorime- 
ter, shown  in  the  last  figure.  This  consists  of  a  perfo- 
rated receptacle,  A,  which  is  surrounded  by  the  ice  to  be 
melted  and  covered  by  a  movable  lid.  This  receptacle 
and  the  ice  chamber  around  it  are  drained  by  the  faucet,  f. 
Outside  of  this  ice  chamber  is  a  second  one  drained  by  the 
faucet,  g,  the  object  of  which  is  to  protect  the  inner  ice 
chamber  from  atmospheric  heat.  Both  chambers  are  cov- 
ered by  a  lid  holding  ice. 

Suppose  now  that  we  wish  to  find  the  specific  heat  of 
iron  and  that  we  take  for  experiment  a  ball  of  the  pure 
metal  weighing  250  Gm.  We  bring  this  to  a  temperature 
of  100°  C,  exactly,  and  then,  without  giving  it  an  oppor- 
tunity to  cool  in  the  air,  drop  it  immediately  into  the  ves- 
sel, A,  and  replace  the  lid.  The  ice  around  A  has  the 
temperature  0°.  Some  of  it  melts,  and  the  water  formed 
is  after  a  time  drawn  off  through  f,  and  weighed  or  meas- 
ured. Assume  that  the  water  collected  weighs  38  grams. 
This  is  necessarily  produced  by  the  melting  of  38  grams 
of  ice;  but  to  melt  one  gram  of  ice  requires  the  addition 
of  79.5  units  of  heat.  Therefore, 

79.5X38  —  3021  units 

^are  given  to  melt  the  38  grams  of  ice.  These  units  are 
furnished  by  the  cooling  of  250  grams  of  iron  through 
100°,  because  the  end  temperature  of  the  metal  must  be 
0°,  with  a  great  excess  of  ice  in  the  apparatus.  It  follows, 
therefore,  that  one  gram  of  iron  in  cooling  down  one  de- 
gree must  give  out  an  amount  of  heat  found  by  this 
division: 

3021 


That  is,  each  gram  of  iron  gives  out  0.12  unit  of  heat  in 
cooling  one  degree,  C.  To  warm  one  gram  of  iron  one 
degree  would  require  the  addition  of  the  same  amount  of 
heat,  and  this,  by  definition,  is  the  specific  heat. 


270  GENERAL  CHEMISTRY. 

The  method  here  "briefly  described  was  suggested  by 
Black  and  improved  by  Lavoisier  and  Laplace,  who  intro- 
duced the  form  of  calorimeter  figured.  The  greatest 
accuracy  is  not  possible  with  this  apparatus,  and  besides 
relatively  large  weights  of  substance  are  required  in  the 
experiment,  but  it  serves  well  to  demonstrate  the  principle. 
Bunsen  later  constructed  an  ice  calorimeter,  which  is  used 
in  very  exact  investigations  with  even  small  amounts  of 
substance.  In  this  apparatus  the  amount  of  melted  ice  is 
determined  by  the  decrease  of  volume  which  follows  when 
it  is  converted  into  water. 

In  another  general  method  what  is  known  as  a  water 
calorimeter  is  used.  In  this  the  elevation  of  temperature 
in  a  given  weight  of  water  is  measured,  when  to  it  a  given 
weight  of  the  substance  at  a  given  temperature  is  added. 
Let  a  represent  the  weight  of  water  in  the  calorimeter,  b 
the  weight  of  substance  added  to  it,  t'  the  original  temper- 
ature of  the  water,  t"  the  temperature  after  mixing,  t"' 
the  temperature  of  the  substance  before  adding  it  to  the 
water,  then  the  specific  heat,  x,  of  the  substance  is  given 
by  the  formula: 

_a(t"—  t') 


This  simple  equation  does  not  include  a  correction 
which  in  practice  must  be  made  for  the  amount  of  heat 
absorbed  by  the  calorimeter  itself.  Other  general  methods 
employed  in  specific  heat  determinations  are  essentially 
modifications  of  the  above. 


CHAPTER  XIII. 


THE  ALKALI   METALS:    LITHIUM,    SODIUM,  POTAS- 
SIUM,   RUBIDIUM    AND    CAESIUM.— 
AMMONIUM  COMPOUNDS. 


GENERAL  CHARACTERISTICS. 

THESE  metals  form  a  natural  group  or  family  as  already 
shown,  and  the  resemblances  extend  through  most  of 
the  compounds.  The  salts  of  ammonium  are  very  similar  to 
the  alkali  metal  salts  in  many  respects  and  therefore  may 
be  conveniently  and  properly  described  with  them.  The 
physical  properties  of  the  metals  are  shown  in  the  follow- 
ing table: 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 

VOLUME. 

MELTING 
POINT. 

BOILING 
POINT. 

Lithium  . 

7  03 

59 

11  9 

180° 

Sodium  

23  05 

97 

23  7 

97 

740° 

Potassium  
Rubidium  
Caesium  

39.11 
85.43 
132.89 

.87 
1.52 
1  88 

44.8 
50.2 
70  6 

60 
38.5 
26.5 

720 

The  metals  resemble  each  other  closely  in  their  behav- 
ior toward  oxygen  and  water.  They  become  coated  with 
a  film  of  oxide  in  dry  air,  and  when  heated  burn  readily. 
In  water  they  dissolve,  forming  strongly  alkaline  hydrox- 
ides. Very  few  of  the  salts  formed  by  these  metals  are  in- 
soluble in  water,  and  these  few  are  not  important.  The 
carbonate  and  phosphate  of  lithium  are  but  slightly  solu- 


272  GENERAL  CHEMISTRY. 

bleand  in  this  respect  bear  some  resemblance  to   the  cor- 
responding magnesium  compounds. 


LITHIUM. 

Occurrence.  Very  widely  distributed  in  nature,  be- 
ing found  in  traces  in  all  mineral  waters,  as  chloride,  car- 
bonate or  sulphate.  It  is  found  in  several  minerals,  in 
many  vegetable  substances  and  has  been  detected  in  the 
blood.  Yet  it  must  be  considered  as  one  of  the  rare 
metals. 

History.  The  metal  was  discovered  in  1817  by  Arf- 
vedson,  was  isolated  by  Davy,  and  first  obtained  in  pure 
form  by  Bunsen  and  Matthiessen  in  1855  by  electrol- 
ysis of  the  fused  chloride. 

Preparation.  The  metal  is  still  isolated  by  the  proc- 
ess last  given. 

Uses.     The  metal  has  no  practical  uses. 

Salts  of  Lithium.  Many  compounds  of  the  metal  are 
known,  but  they  are  not  abundant  enough  to  be  impor- 
tant. Lithium  chloride  is  extremely  soluble  in  water. 
Lithium  phosphate  requires  about  2,500  parts  of  water 
for  its  solution.  Compounds  of  lithium  are  best  recog- 
nized by  the  color  they  impart  to  the  Bunsen  flame  and 
by  the  characteristic  red  line  in  its  spectrum. 


SODIUM. 

Occurrence.  This  is  one  of  the  most  widely  distrib- 
uted of  the  elements.  It  is  found  in  largest  quantity  as 
the  chloride  in  rock  salt,  sea  salt  and  in  many  mineral 
springs;  as  nitrate  in  Chili  saltpeter;  as  sulphate  and  car- 
bonate in  large  deposits  in  several  parts  of  the  world  and 
in  many  other  compounds. 

History.  The  metal  was  first  isolated  by  Davy  in 
1807  by  electrolysis  of  fused  soda.  It  was  later  made  by 


GENERAL  CHEMISTRY.  273 

heating  iron  with  caustic  soda;  and  in  greater  quantity  by 
heating  a  mixture  of  sodium  carbonate  and  charcoal. 

Manufacture.  Until  recently  sodium  has  been  used 
in  large  quantities  for  the  production  of  aluminum,  and  this 
employment  greatly  stimulated  the  industry.  Sodium  is 
practically  produced,  on  the  large  scale  by  the  reduction  of 
the  carbonate  or  hydroxide  by  some  form  of  carbon,  but  the 
general  process  has  undergone  several  modifications.  By 
one  method  a  mixture  is  made  of  8  parts  of  dry  sodium 
carbonate,  3  parts  of  powdered  coal  or  charcoal  and  1  part 
of  chalk.  This  is  mixed  to  a  stiff  paste  with  oil  or  tar, 
charged  into  an  iron  retort,  and  distilled.  The  iron  retort 
is  connected  by  a  pipe  with  an  iron  box  which  is  kept  cold 
outside  the  furnace.  By  the  decomposition  of  the  chalk 
CO2  is  liberated,  which  helps  to  carry  along  the  sodium 
vapors.  These  condense  in  the  iron  box  which  is  emptied 
at  the  end  of  the  distillation,  the  sodium  being  poured  into 
petroleum.  The  reaction  illustrating  this  reduction  is 
this: 


In  another  process  the  reducing  agent  is  a  crude  iron 
carbide  obtained  by  heating  a  mixture  of  finely  divided 
iron  and  tar  and  cokin'g  the  same  away  from  the  air.  This 
yields  a  body  of  the  formula  FeC2.  This  carbide  is  mixed 
with  caustic  soda  in  the  proportion  of  10  parts  of  the  latter 
to  7.5  of  the  former,  and  the  mixture  is  distilled  as  before 
in  a  retort.  The  reaction  takes  place  at  a  relatively  low 
temperature  and  according  to  the  equation: 


The  escaping  gases  assist  in  carrying  the  sodium  vapor 
from  the  retort. 

Two  electrolytic  processes  are  also  in  use  for  the  separa- 
tion of  sodium.  In  one  the  hydroxide,  in  molten  condition, 
is  decomposed  at  a  temperature  but  little  above  its  fusing 
point,  which  permits  the  easy  removal  of  the  liberated 
metal.  In  another  process  the  chloride  is  mixed  with 


274  GENERAL  CHEMISTRY. 

potassium  or  calcium  chloride,  fused  at  alow  red  heat,  and 
subjected  to  electrolysis. 

Properties.  The  specific  gravity,  melting  and  boiling 
points  of  sodium  have  been  given  above.  The  metal  is 
silver  white  but  tarnishes  readily  in  the  air.  It  decomposes 
water  and  other  hydroxyl  compounds  readily  with  libera- 
tion of  hydrogen.  It  is  an  interesting  fact  that  with  dry 
chlorine  gas  or  pure  liquid  chlorine  or  bromine  it  is  prac- 
tically inert.  In  dry  hydrochloric  acid  gas  it  remains 
unattacked  for  weeks.  It  will  be  seen  in  what  follows  that 
the  same  behavior  is  observed  with  other  metals. 

Uses.  Sodium  has  been  employed  for  the  production 
of  aluminum  by  the  decomposition  of  the  chloride  at  a  high 
temperature. 

AiCl8+3Na=Al+aNaCl. 

This  process  has  been  practically  displaced  by  the 
cheaper  electrolytic  method.  The  separation  of  magne- 
sium from  the  chloride  is  accomplished  in  the  same 
manner: 

MgCl8+2Na  =  Mg+2NaCl. 

It  is  employed  in  the  production  of  sodium  amalgam 
which  is  used  for  several  purposes,  also  in  making  pure 
sodium  hydroxide  for  some  laboratory  purposes  and  in 
various  other  ways. 

Sodium  Oxides.  Several  of  these  bodies  have  been 
described  and  one  of  them  has  recently  become  important. 
A  mixture  of  Na2O  and  Na2O2is  formed  by  burning 
sodium  in  dry  air  or  oxygen.  The  oxide  Na2O2,  known 
in  commerce  as  the  dioxide,  peroxide  or  superoxide, 
is  made  by  passing  pure  dry  air,  CO2  free,  over  metal- 
lic sodium  heated  in  an  iron  pipe  placed  in  a  furnace. 
The  sodium  burns  slowly  and  a  white  mass  results.  This 
is  used  in  analytical  chemistry  and  technology  as  an  oxi- 
dizing agent,  being  valuable  because  of  the  readiness  with 
which  it  gives  up  oxygen. 

Sodium    Hydroxide.     This   important    compound  is 


GENERAL  CHEMISTRY.  275 

known  as  caustic  soda  and  is  made  in  enormous  quantities 
for  use  in  several  industries.  It  is  usually  prepared  from 
the  carbonate  by  boiling  a  solution  of  the  latter  with 
slaked  lime  : 

Na2C03-fCa02H3=2NaOH+CaC03. 
On  a  smaller  scale  it  is  made  by  action  of  water: 

Na-f-H2O  =  NaOH-f-H. 
The  following  experiment   illustrates  this  : 

Ex.  153.  Cut  a  gram  or  more  of  sodium  into  small  bits,  which 
throw  into  a  beaker  containing  about  100  Cc.  of  water,  one  at  a  time, 
waiting  until  the  action  ceases  in  each  case  before  adding  another  piece. 
Cover  the  beaker  with  a  glass  plate  during  the  reaction.  When  this  is 
complete,  remove  the  glass  and  evaporate  the  solution  to  dryness,  be- 
ginning in  the  beaker  and  finishing  in  a  small  porcelain  dish.  Observe 
the  color  of  the  solid  substance  obtained,  and  its  action  on  the  skin. 
Now  add  about  50  Cc.  of  water  to  the  solid  to  dissolve  it.  Divide  the 
solution  into  four  portions.  To  one  add  a  few  drops  of  solution  of 
phenol-phthalein  in  alcohol ;  to  the  second  some  solution  of  litmus,  to 
which  a  drop  or  two  of  acid  had  been  added  ;  to  the  third  some  aqueous 
solution  of  methyl  orange  to  which  a  trace  of  acid  had  been  added  ;  and 
to  the  fourth  some  solution  of  cochineal  in  weak  alcohol,  made  acid  with 
a  few  drops  of  hydrochloric  acid.  Observe  the  color  of  these  solutions 
before  they  are  added  to  the  sodium  hydroxide. 

In  each  case  a  marked  change  in  shade  follows  the 
mixing.  The  litmus,  phenol-phthalein,  methyl  orange  and 
cochineal  are  organic  substances,  which  have  one  color  in 
acid  solutions  and  another  in  alkaline  solutions.  We  have 
already  had  illustrations  of  this  behavior,  and  we  shall  find 
others  later.  Our  sodium  hydroxide  is  one  of  the  most 
characteristic  of  alkaline  substances,  and  the  colors  pro- 
duced above  are  essentially  the  same  as  are  produced  by 
other  alkalies. 

The  production  from  the  carbonate  is  illustrated  by  the 
next  experiment. 

Ex.  154.  Dissolve  50  Gm.  of  common  crystallized  sodium  car- 
bonate in  250  Cc.  of  water  in  a  porcelain  dish,  and  add  milk  of  lime 
made  by  mixing  20  Gm.  of  slaked  lime  with  75  Cc.  of  water.  The  milk 
of  lime  is  to  be  added  slowly  to  the  boiling  solution  of  the  carbonate, 
the  mixture  being  well  stirred  during  the  addition.  After  the  lime  has 
been  all  added,  boil  the  mixture  about  ten  minutes,  keeping  it  well 


276  GENERAL  CHEMISTRY. 

stirred.  Filter  the  liquid  portion  through  a  plug  of  asbestos  in  the  apex 
of  a  funnel  and  then  pour  on  the  sediment  in  the  dish.  When  the  liquid 
has  drained  from  this,  add  a  little  water  to  the  funnel  to  wash  out  more 
of  the  product.  Test  a  portion  of  the  filtrate  with  hydrochloric  acid. 
It  should  give  no  effervescence  from  presence  of  unchanged  carbonate. 
Evaporate  the  remainder  in  a  porcelain  dish  and  observe  that  a  white 
residue  of  soda  is  left.  During  the  evaporation  this  takes  up  some  car- 
bon dioxide  from  the  air  and  becomes  partly  carbonate  again. 

On  the  large  scale  this  reaction  is  carried  out  in  iron 
boilers.  The  hydroxide  solution,  separated  from  the  sedi- 
ment, is  boiled  down  in  kettles,  and  when  ready  to  solidify 
is  poured  into  iron  drums  for  shipment  in  large  lots  or  into 
molds  for  formation  of  smaller  lumps.  Some  is  cast  in 
stick  form  for  laboratory  use.  In  chemical  industry  the 
hydroxide  is  employed  in  making  soap  by  the  saponifica- 
tion  of  fats;  also  in  the  production  of  many  organic  color- 
ing matters,  and  for  minor  purposes.  The  commercial 
article  usually  contains  a  little  water. 

Several  other  methods  of  producing  soda  have  been 
tried.  The  following  experiment  will  illustrate  the  prin- 
ciple involved  in  a  method  which  promises  at  the  present 
time  to  make  great  changes  in  the  alkali  industry,  and 
which  is  already  in  successful  operation  on  the  large 
scale. 

Ex.  155.  Nearly  fill  a  large  U-tube  with  a  strong  solution  of 
sodium  chloride  Into  each  limb  of  the  tube  dip  thin  pieces  of  platinum 
foil,  about  1  Cm.  wide  and  5  Cm.  long.  These  pieces  of  platinum  are 
fastened  to  platinum  wires,  which  in  turn  are  attached  to  copper  wires 
leading  from  the  poles  of  a  galvanic  battery  of  five  or  six  Daniell  or 
equivalent  elements.  Just  before  making  the  last  attachments  pour  a 
few  drops  of  solution  of  phenol-phthalein  into  the  limb  connected  with 
the  zinc  pole  of  the  battery,  and  some  indigo  solution  into  the  other 
limb  of  the  U-tube.  Now  complete  the  connection  with  the  battery, 
and  observe  that  the  salt  solution  in  the  side  with  the  phenol-phthalein 
soon  turns  bright  red,  showing  formation  of  alkali,  while  the  indigo 
solution  in  the  other  limb  is  gradually  bleached  by  chlorine  liberated, 
the  odor  of  which  can  be  readily  detected.  The  electric  current  pass- 
ing through  the  solution  decomposes  the  salt,  separating  the  sodium 
from  the  chlorine.  The  latter  is  liberated  on  one  plate,  or  electrode, 
while  the  former  decomposes  water  at  the  other,  producing  caustic 
soda.  On  the  large  scale  it  has  been  found  possible  to  separate  and 
save  the  soda  so  formed,  powerful  currents  from  dynamos  being 
employed. 

Sodium    Carbonate.     This    substance,     mixed   with 


GENERAL  CHEMISTRY. 


277 


sodium  sulphate,  is  found  in  great  deposits  in  Wyoming 
and  other  parts  of  the  United  States.  As  no  easy  method 
of  purification  has  yet  been  discovered,  this  crude  soda 
has  been  utilized  to  a  limited  extent  only. 

Leblanc  Process.  The  commercial  carbonate  is  pro- 
duced by  two  essentially  different  processes  at  the  present 
time.  The  older  of  these  is  known  as  the  Leblanc 
process  and  depends  on  these  reactions.  Common  salt  is 
converted  into  sulphate  by  the  action  of  strong  sulphuric 
acid: 

2NaCl+H8S04  =  Na8S 


The  sulphate  is  roasted  with  a  mixture  of  coal  dust  and 
limestone.     The  first  reduces  sulphate  to  sulphide  : 


The  second  then  converts  sulphide  into  carbonate  : 
Na2S+CaCO3=Na2CO3+CaS. 

The  first  one  of  these  reactions  is  conducted  inarever- 
beratory  furnace  so  constructed  that  the  hydrochloric  acid 
maybe  led  off  into  Woulfe  bottles  and  condensed,  or,  more 
commonly  into  coke  towers  and  there  condensed  by  water 
flowing  down  over  the  coke.  The  sulphate  is  thoroughly 
roasted  to  expel  all  acid  and  in  this  form  is  known  as  salt 
cake.  This  is  broken  into  lumps  and  introduced  into  a 
second  furnace  with  the  coal  and  limestone.  This  may  be 
a  reverberatory  furnace  also,  and  is  commonly  called  a 
black  ash  furnace,  because  the  three  substances  on 
thorough  roasting  yield  a  black  mass.  At  the  end  of  the 
process  this  is  raked  out,  cooled  and  leached  with  water. 
The  calcium  sulphide  is  not  soluble  in  the  alkaline  liquid 
while  the  carbonate  dissolves  readily,  with  several  by- 
products. The  leach  liquor  is  evaporated  to  crystallize  the 
carbonate  which  is  deposited  in  the  form  of  sal  soda, 
Na2CO3.10H2O.  A  large  part  of  this  product  is  calcined 
to  drive  off  the  water.  The  residue  is  known  as  soda  ash 
and  is  largely  used  in  the  production  of  common  glass.  A 
part  of  the  sal  soda  is  purified  by  recrystallization  and  is 


278  GENERAL  CHEMISTRY. 

used  in  making  many  pharmaceutical  products.  It  is  also 
used,  as  shown  above,  in  making  caustic  soda.  Some  of 
the  important  properties  of  the  sal  soda  may  be  learned 
from  the  following  experiment  : 

Ex.  156.  Let  the  student  make  some  experiments  with  a  large 
crystal  of  sal  soda.  Break  it  into  small  pieces.  Heat  some  of  these  in 
a  porcelain  dish  and  observe  that  they  melt  to  a  liquid  mass  which  gives 
off  water  by  continued  heating,  leaving  finally  a  white  powder.  Dis- 
solve some  of  the  fragments  of  the  crystal  soda  in  water.  Notice  the 
taste  of  the  solution  by  bringing  a  small  drop  on  a  glass  rod  to  the 
tongue.  Test  the  reaction  of  the  solution  with  litmus  and  phenol- 
phthalein.  Pour  some  dilute  hydrochloric  acid  into  a  little  of  the  solu- 
tion in  a  beaker.  A  gas  is  given  off.  What  is  it  ?  Continue  the  addi- 
tion of  the  dilute  hydrochloric  acid  until  no  more  gas  escapes  on  shaking, 
then  evaporate  the  solution  to  complete  dryness  in  a  porcelain  dish. 
Taste  the  residue.  What  is  it  ?  Into  some  more  of  the  solution  of  the 
crystals  to  which  a  little  phenol-phthalein  has  been  added,  pass  carbon 
dioxide  gas  from  a  generator  until  the  color  disappears.  Sodium  bicar- 
bonate is  here  formed. 

Solvay  Process.  The  largest  part  of  the  carbonate  of 
commerce  is  now  made  by  a  process  known  as  the  Solvay 
process  or  ammonia  process.  This  is  the  result  of  attempts 
to  decompose  salt  directly  without  the  aid  of  the  sulphuric 
acid,  and  depends  on  the  reaction: 


=  NH4Cl+HNaCOs. 

Strong  brine  is  saturated  with  ammonia  gas  and  then 
treated  with  carbon  dioxide  in  tall  tanks  under  pressure. 
As  the  ammoniacal  brine  becomes  saturated  a  fine  granular 
precipitate  of  sodium  bicarbonate  forms  and  gradually 
sinks  to  the  bottom  of  the  reaction  tanks.  This  continues 
until  a  large  part  of  the  salt  is  decomposed.  The  bicar- 
bonate is  now  removed,  washed  with  cold  water  saturated 
with  the  same  salt,  dried  and  calcined  to  produce  soda  ash. 
The  CO2  given  off  is  utilized  for  saturation  again.  The 
success  of  the  process  depends  on  the  fact  that  from  the 
NH4C1  left  in  the  mother  liquor  the  whole  of  the  ammonia 
may  be  recovered,  by  boiling  with  slaked  lime,  and  used 
over  again  with  fresh  brine. 

Sodium  Bicarbonate.     As  shown  at  the  conclusion  of 


GENERAL  CHEMISTRY.  279 

the  last  experiment,  sodium  bicarbonate  is  formed  when 
carbon  dioxide  gas  is  led  into  a  solution  of  the  carbonate, 
or  when  the  gas  is  brought  in  contact  with  lumps  of  par- 
tially effloresced  sal  soda. 

C02+Na2C03+H20  =  2HNaC03. 

As  the  behavior  with  the  phenol-phthalein  shows,  this 
is  not  an  alkali  substance.  It  is  made  in  great  quantities 
for  use  in  the  manufacture  of  baking  powders,  effervescing 
mixtures  and  other  medicinal  products. 

Ex.  157.  Let  the  student  mix  a  small  amount  of  finely  powdered 
"  cream  of  tartar  "  (an  acid  substance),  with  some  sodium  bicarbonate 
on  a  sheet  of  paper.  No  change  is  apparent.  Then  throw  the  mixture 
into  a  beaker  and  add  some  water.  A  lively  effervescence  follows.  This 
is  from  the  escape  of  the  CO2,  set  free  by  the  action  of  the  acid  sub- 
stance on  the  bicarbonate  in  presence  of  water. 

Sodium  Chloride,  or  common  salt,  is  the  most  abun- 
dant of  the  sodium  compounds.  It  is  obtained  from  de- 
posits of  rock  salt  found  in  many  parts  of  the  world,  from 
brine  springs,  and  from  sea  water.  In  very  cold  latitudes 
this  is  concentrated  to  a  point  where  boiling  is  economic- 
ally possible  by  freezing  the  water  pumped  up  into  large 
shallow  basins.  In  freezing  little  of  the  salt  separates  with 
the  ice.  The  operation  is  repeated  several  times,  until 
the  greater  part  of  the  water  is  removed.  In  tropical  or 
semitropical  countries  water  is  pumped  from  the  ocean  or 
allowed  to  flow  at  high  tide  into  basins,  where  evaporation 
takes  place  by  the  sun's  heat.  The  brine  of  many  springs 
is  strong  enough  in  salt  to  pay  for  direct  concentration 
over  fire.  After  separation  of  most  of  the  salt  the  mother 
liquor  is  used  for  the  manufacture  of  bromides,  which  in 
many  places  is  an  important  industry.  As  shown  above, 
salt  is  used  in  making  other  important  sodium  compounds. 
It  dissolves  in  somewhat  less  than  three  parts  of  water  at 
the  ordinary  temperature. 

Sodium  iodide,  Nal,  and  sodium  bromide,  NaBr,  are 
analogous  compounds,  used  mainly  in  medicine. 

Sodium  Sulphite  is  a  soluble  salt,  made  by  passing 


280  GENERAL  CHEMISTRY. 

SO2  into  a  solution  of  the  carbonate  to  saturation,  after 
which  an  equal  amount  of  carbonate  is  added,  and  the 
whole  allowed  to  stand  to  crystallize  as  Na2SO3.7H2O. 
The  salt  is  used  for  several  purposes,  and  very  commonly 
as  a  constituent  of  developing  solutions  in  photography, 
on  account  of  its  reducing  properties.  The  acid  sulphite, 
HNaSO3,  is  also  used  in  photography. 

Sodium  Sulphate.  This  is  a  very  common  substance 
made  by  the  action  of  sulphuric  acid  on  common  salt.  In 
the  crystalline  form,  Na2SO4.10H2O,  it  is  known  as  Glau- 
ber's salts.  A  partially  effloresced  product  is  found  in 
large  deposits  in  the  western  part  of  the  United  States. 
The  dried  sulphate  is  often  used  in  the  place  of  the  car- 
bonate in  making  glass,  but  a  higher  heat  is  required  in 
the  fusion. 

Sodium  Thiosulphate,  Na2S2O3.5H2O,  is  made  in 
large  quantities  for  use  in  photography  as  a  fixing  agent, 
and  to  remove  excess  of  chlorine  after  bleaching  in  the 
paper  industry.  Its  action  in  photography  will  be 
explained  later. 

Sodium  Nitrate.  This  important  compound  is  found 
in  large  deposits  along  the  western  coast  of  South  America 
and  is  known  as  Chili  saltpeter.  In  crude  form  it  jcomes 
into  commerce  for  use  as  a  fertilizer.  After  refining  it  is 
employed  in  making  nitric  acid  and  in  making  potassium 
nitrate,  to  be  presently  explained.  Large  quantities  are 
also  used  in  the  manufacture  of  gunpowder. 

Among  other  important  compounds  of  sodium  there  are 
borax,  Na2B4O7.10H2O,  and  sodium  silicate  or  water-glass, 
Na2SiO3,  described  in  the  eighth  chapter  and  the  phos- 
phates, described  in  the  ninth  chapter.  Sodium  dichromate, 
Na2Cr2O7,  is  an  important  salt  which  will  be  referred  to 
later. 

Recognition.  Sodium  compounds  are  easily  recognized 
by  the  intense  yellow  color  they  impart  to  the  flame  of  the 
Bunsen  burner  and  by  the  bright  yellow  line  seen  with  the 
spectroscope,  to  be  explained  in  Chapter  XV. 


GENERAL  CHEMISTRY.  281 

POTASSIUM. 

Occurrence.  Found  very  widely  distributed  in  rocks 
and  soils.  As  potassium  carbonate  it  occurs  in  the  ashes 
ot  plants.  Potassium  chloride,  KC1,  occurs  as  silvine  in 
the  famous  salt  deposits  of  Stassfurt,  Germany,  along 
with  carnallite,  KC1.  MgCl2.6H2O.  A  large  part  of  the 
potassium  compounds  of  commerce  is  made  from  these 
minerals.  Potassium  nitrate  appears  as  an  efflores- 
cence on  many  soils  in  hot  climates,  and  at  one  time 
this  was  a  very  important  source  of  the  crude  niter  used 
in  Europe. 

History.  What  was  said  about  sodium  applies  to 
potassium.  It  was  isolated  in  1807  by  Davy  from  potassa, 
in  which  the  presence  of  a  metal  was  not  suspected.  The 
compounds  we  call  potassa  and  soda  were  supposed  to 
be  simple  substances  by  most  chemists,  although  Lavoisier 
had  suggested  that  they  must  contain  oxygen. 

Manufacture.  The  reactions  given  for  sodium  apply 
here  as  well.  In  the  early  aluminum  industry  potassium 
was  employed  as  the  reducing  agent,  but  because  of  its 
much  greater  atomic  weight  was  later  displaced  by  sodium 
which  could  be  produced  even  more  readily.  At  the  pres- 
ent time  potassium  is  not  manufactured  in  large  quan- 
tities. 

Properties.  A  soft  white  metal  which  tarnishes  read- 
ily, burns  when  heated  in  the  air  and  decomposes  water 
with  liberation  of  hydrogen.  In  general  its  behavior  is 
much  like  that  of  sodium. 

Uses.  It  is  employed  sometimes  in  scientific  investi- 
gations and  in  the  preparation  of  the  pure  hydroxide  for 
experimental  purposes. 

The  compounds  of  potassium  closely  resemble  those  of 
sodium. 

Oxides.  Two  oxides,  K2O  and  K2O4,  are  well  known 
but  not  technically  important.  Several  other  oxides  have 
been  described. 


282  GENERAL  CHEMISTRY. 

Potassium  Hydroxide.     The  pure  substance  may  be 
made  by  the  action  of  potassium  on  water, 


but  is  commonly  made  by  the  decomposition  of  potassium 
carbonate  by  slaked  lime  as  described  under  sodium: 

K2CO3+CaO2H2=,CaCO3-j-2KOH. 

The  CaCO3  we  are  already  familiar  with  as  calcium 
carbonate,  identical  in  composition  with  marble  and  pure 
limestone.  It  is  insoluble  in  water,  and  at  the  end  of  the 
decomposition  remains  in  the  sediment  with  the  excess 
of  CaO2H2  taken. 

Potassium  hydroxide  is  employed  for  many  purposes  in 
the  arts,  especially  in  the  manufacture  of  soft  soap  and 
certain  organic  coloring  substances.  For  laboratory  and 
pharmaceutical  uses  it  is  sold  in  the  form  of  sticks,  for 
convenience  in  handling.  This  product  is  not  pure  KOH, 
but  contains  always  10  to  15  per  cent  of  water,  with  some 
foreign  substances. 

Potassium  Carbonate.  This  very  important  com- 
pound is  produced  on  the  large  scale  in  several  ways,  but 
only  one  method  of  manufacture  will  be  illustrated  here. 
Wood  ashes  contain  notable  quantities  of  the  carbonate 
which  can  be  extracted  with  water.  The  solution  obtained 
yields  on  evaporation  a  crude  carbonate. 

Ex.  158.  Pour  lukewarm  water  over  several  hundred  grams  of  fine 
wood  ashes  in  a  large  beaker.  Enough  water  should  be  taken  to  make  a 
thin  mixture  easily  filtered.  After  stirring  the  mixture  thoroughly  filter 
it  through  a  large  filter,  or  better,  through  a  plug  of  asbestos  in  the  neck 
of  a  funnel.  The  filtrate  will  have  a  brownish  color.  Evaporate  it  to 
dryness  in  a  porcelain  dish.  The  residue  obtained  contains  a  large  pro- 
portion of  potassium  carbonate,  with  other  soluble  salts  from  the  ash, 
as  potassium  sulphate,  potassium  chloride,  potassium  silicate  and  other 
products  in  smaller  amount.  Test  the  alkalinity  of  the  substance  by 
means  of  litmus  or  phenol-phthalein,  and  try  the  behavior  with  acids. 

The  crude  potash  may  be  partially  purified  by  heating 
very  strongly  to  destroy  organic  matters  present.  The 
dry  product  is  dissolved  in  twice  its  weight  of  water,  fil- 


GENERAL  CHEMISTRY.  283 

tered,  the  filtrate  evaporated  to  one-third  its  volume  and 
allowed  to  stand.  Most  of  the  sulphate  and  some  other 
salts  settle  cut.  The  remaining  liquid  is  evaporated  until 
the  carbonate  begins  to  crystallize,  and  is  then  allowed  to 
cool  and  deposit  most  of  it.  This  product  is  collected, 
dried  and  strongly  heated  to  drive  off  more  water  and 
yields  then  what  is  known  as  pearl  ash  or  refined  potash. 
Much  of  the  carbonate  of  commerce  is  now  made  by  the 
Leblanc  and  Solvay  processes  as  described  under  sodium. 

When  perfectly  dry  it  is  used  to  absorb  water  from  cer- 
tain liquids,  as  commercial  alcohol,  to  dry  gases,  and  for 
other  purposes  depending  on  its  power  of  attracting  mois- 
ture. This  power  can  be  tested  by  exposing  some  of  the 
dry  product  to  the  air  for  several  days. 

The  carbonate  is  used  in  the  manufacture  of  the  hydrox- 
ide, as  stated,  and  also  in  making  certain  kinds  of  glass. 
Hard  glass  is  essentially  a  combination  of  the  silicates  of 
calcium  and  potassium  obtained  by  melting  together  sand, 
lime  and  calcined  potash. 

Potassium  Chloride.  The  native  mineral  is  found  in 
nearly  pure  condition,  and  by  recrystallization  is  made 
ready  for  laboratory  use.  It  serves  as  the  starting  point  in 
making  other  potassium  salts,  just  as  common  salt  serves 
for  the  sodium  compounds. 

Potassium  Iodide  and  Bromide  are  important  sub- 
stances used  in  medicine.  They  are  produced  from  free 
iodine  and  bromine  by  several  reactions,  which  cannot  be 
described  here. 

Potassium  Sulphate  is  made  by  the  action  of  sul- 
phuric acid  on  the  chloride.  It  is  used  for  several  pur- 
poses in  the  laboratory,  and  on  the  large  scale  serves  for 
the  preparation  of  potassium  carbonate  by  the  Leblanc 
process. 

Potassium  Nitrate.  The  important  properties  of  this 
substance  were  explained  when  nitric  acid  was  studied.  It 
is  very  soluble  in  water,  easily  melted,  and  is  readily  de- 
composed by  sulphuric  acid,  yielding  nitric  acid  and  potas- 


284  GENERAL  CHEMISTRY. 

sium  sulphate.  When  mixed  with  sulphur  and  charcoal  it 
gives  up  oxygen  readily  under  certain  conditions  and  often 
with  explosive  suddenness.  Hence  its  use  in  gunpowder 
and  fireworks  mixtures. 

It  is  manufactured  in  great  quantities  and  generally  by 
decomposition  between  potassium  chloride  and  sodium 
nitrate: 

KCl+NaNO3  =  KNO3-fNaCl. 

The  reaction  depends  on  the  facts  that  potassium 
nitrate  is  very  soluble  in  hot  water  and  much  more  soluble 
than  sodium  chloride,  and  further  that  the  latter  is 
not  much  more  soluble  in  hot  water  than  in  cold. 
The  manufacture  is  carried  out  in  various  ways; 
sometimes  by  mixing  the  chloride  and  nitrate  in  molecular 
proportions,  that  is  74.6  parts  of  KC1  to  85  parts  of  NaNO3, 
and  adding  this  mixture  to  boiling  water  as  long  as  it  all 
dissolves.  When  a  condition  of  saturation  is  reached  the 
liquid  is  allowed  to  cool.  A  large  amount  of  KNO3  crys- 
tallizes out,  while  nearly  all  the  NaCl  produced  stays  in 
solution.  The  saltpeter  is  then  easily  purified  by  recrys- 
tallization  from  water.  A  more  common  method  is  to  add  the 
molecular  mixture  to  boiling  water,  usually  the  water  from 
which  saltpeter  has  been  recrystallized  in  previous  opera- 
tions, until  a  specific  gravity  of  1.5  is  reached.  Sodium 
chloride  is  precipitated  in  this  way  and  the  supernatant 
liquid  drawn  off  deposits  the  potassium  nitrate  on 
cooling. 

Potassium  Chlorate.  This  salt  is  usually  made  by  a 
series  of  reactions  which  are  somewhat  complicated.  When 
chlorine  gas  is  passed  into  a  strong,  hot  solution  of  potas- 
sium hydroxide  this  change  takes  place: 

3C12+6KOH  =  5KC1+KC1O3+3H2O, 

and  in  this  manner  the  salt  was  formerly  made,  being  easily 
separated  from  the  KC1  by  crystallization.  At  present  the 
much  cheaper  substance,  milk  of  lime,  CaO2H2,  is  con- 
verted in  the  same  way  into  calcium  chlorate,  Ca^lO^. 
The  solution  is  concentrated  and  treated  with  KC1  which 


GENERAL  CHEMISTRY.  286 

brings  about  a  double  decomposition,  because  potassium 
chlorate  is  less  soluble  than  calcium  chlorate: 

Ca(ClO3)2+2KCl  =  2KClO3+CaCl2. 

The  chlorate  is  easily  purified  by  crystallization,  as  it  is 
not  very  soluble  in  cold  water. 

This  salt,  like  the  nitrate,  is  valuable  mainly  because  of 
its  power  of  furnishing  oxygen.  Hence  its  use  in  the 
manufacture  of  oxygen  gas,  already  illustrated.  In  the 
manufacture  of  explosive  mixtures  and  of  fireworks  it 
plays  an  important  part.  Some  formulas  are  here  given 
illustrating  the  production  of  these  mixtures  for  colored 
fires: 


WHITE. 

Potassium  nitrate 16  parts. 

Antimony  sulphide  (nat- 
ural)   6      " 

Sulphur 4      " 

Red  lead 5     " 


VIOLET. 

Potassium  nitrate 16  parts. 

Potassium  chlorate 14      " 

Chalk 10      " 

Sulphur 10      " 

Lampblack 0.5  " 


GREEN. 

Potassium  chlorate. ...  3  parts. 

Barium  nitrate 8      " 

Sulphur 3      " 

BLUE. 

Potassium  chlorate 6  parts. 

Ammonia  copper  s  u  1  - 

phate 8      " 

Shellac 1      " 

RED. 

Potassium  chlorate 3  parts. 

Strontium  nitrate 18      " 

Shellac..  6      " 


In  all  cases  the  materials  used  must  be  dry,  and  must 
be  powdered  separately.  The  proper  weights  of  the  differ- 
ent substances  are  then  mixed,  without  rubbing,  best  on  a 
sheet  of  paper,  with  the  hand.  It  will  be  observed  that 
either  the  nitrate  or  chlorate  of  potassium  appears  in  each 
one  of  these  formulas. 

Among  other  important  potassium  salts  the  chromate 
and  dichromate  must  be  referred  to.  More  will  be  said 
about  them  in  the  chapter  on  chromium.  The  perman- 
ganate will  also  be  described  later. 

Recognition.  Potassium  compounds  are  recognized 
by  the  violet  color  they  impart  to  the  Bunsen  flame,  by  the 
bright  red  line  they  exhibit  in  the  spectroscope  and  by  a 


286  GENERAL  CHEMISTRY. 

yellow  precipitate  formed   under  certain  conditions  with 
platinum  chloride. 

RUBIDIUM  AND  CESIUM. 

These  are  two  rare  metals  found  in  a  number  of  miner- 
als, but  not  in  large  amount.  They  are  found  in  traces  in 
many  mineral  waters  and  it  was  in  the  examination  of  such 
waters  by  the  spectroscope  that  they  were  discovered  in 
1860  and  1861  by  Bunsen.  Caesium  is  characterized  by 
showing  two  bright  blue  lines  and  rubidium  by  a  number 
of  lines,  among  them  two  in  the  deep  red  end  of  the 
spectrum. 

AMMONIUM  COMPOUNDS. 

It  has  been  explained  already  that  ammonia,  a  gaseous 
substance  represented  by  the  formula  NH3,  is  obtained 
usually  from  certain  by-products  formed  in  the  manufac- 
ture of  illuminating  gas.  These  by-products  appear  as 
salts  of  ammonium,  NH4,  a  hypothetical  combination  not 
known  in  the  free  state,  but  assumed  to  exist  in  union  with 
other  elements  in  the  form  of  salts.  The  general  behavior 
of  these  salts  has  been  shown  by  experiment.  They  are 
volatile,  some  being  decomposed  at  the  same  time.  They 
are  nearly  all  extremely  soluble  in  water,  and  all  are  decom- 
posed when  boiled  with  strong  alkali  solutions.  We  can 
apply  these  facts  to  the  detection  of  ammonia  in  mixtures. 

The  most  important  of  the  ammonium  salts  are  the 
chloride,  the  nitrate,  the  sulphate  and  the  carbonate.  The 
first  three  are  usually  made  by  direct  combination  of 
ammonia  water  with  the  corresponding  acids.  The  com- 
mercial carbonate  of  ammonium  is  a  mixture  of  the  bicar- 
bonate or  acid  carbonate,  H(NH4)CO3,  with  ammonium 
carbamate,  (NH4)NH2CO?,  and  is  usually  made  by  sub- 
liming a  mixture  of  ammonium  chloride  aad  chalk.  When 
this  is  dissolved  in  water  with  the  addition  of  ammonia  the 
true  normal  carbonate,  (NH4)2CO3,  is  formed. 

By  saturating  ammonia  water  with  hydrogen  sulphide 
gas  a  solution  is  obtained  which  contains  (NH4)  HS.  This 
is  much  used  as  a  laboratory  reagent.  By  adding  to  this 


GENERAL  CHEMISTRY.  287 

solution  an  equal  volume  of  the  ammonia  water  the  sul- 
phide, (NH4)2S,  is  formed.  Both  solutions  have  applica- 
tions in  analytical  chemistry  and  both  decompose  on 
standing. 

Recognition.  Ammonium  compounds  are  readilyrecog- 
nized  by  the  liberation  of  ammonia  when  they  are  boiled 
with  milk  of  lime  or  other  fixed  alkali  solution.  NH3  is 
the  only  common  volatile  alkali.  Among  organic  com- 
pounds there  are  known  a  large  number  of  bodies  called 
amins,  or  compound  ammonias,  which  show  some  of  the 
properties  of  ordinary  ammonia;  but  as  they  are  not  com- 
monly met  with,  they  are  not  liable  to  be  mistaken  for  the 
inorganic  ammonium  salts. 


CHAPTER  XIV. 


THE  COPPER  GROUP:    COPPER,  SILVER  AND  GOLD. 


GENERAL  CHARACTERISTICS. 

THE  METALS  of  this  group  are  not  as  closely  related 
as  are  those  of  the  alkali  group,  yet  they  show  in  their 
physical  properties  and  in  the  properties  of  their  com- 
pounds many  analogies.  As  metals  they  resemble  each 
other  in  their  extreme  malleability,  ductility  and  position 
of  their  melting  points.  In  many  of  their  compounds  they 
are  univalent,  and  in  this  they  resemble  the  alkali  metals. 
Solutions  of  their  compounds  are  very  readily  electrolyzed 
with  precipitation  of  the  pure  metals.  The  following  table 
exhibits  important  physical  properties: 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

BOILING 
POINT. 

Copper.  .  . 

63.60 

8.95 

7.10 

1045° 

White  heat 

Silver  

107.92 

10.55 

10.23 

945° 

Gold  

197.24 

19.30 

10.22 

1035° 

ii 

COPPER. 

Occurrence.  The  metal  is  found  in  large  quantities  in 
the  free  or  native  condition,  especially  in  the  Lake  Supe- 
rior mines.  In  combination  the  common  ores  are  copper 
pyrites,  CuFeS2;  copper  glance  or  vitreous  copper  ore, 
Cu2S;  malachite,  CuCO3.CuO2H2;  red  copper  ore, Cu2O; 
and  gray  copper  or  tetrahedrite,  4Cu2S.Sb2S3.  Many 


GENERAL  CHEMISTRY.  289 

other  ores  of  less  importance  are  known,  and  these  cited 
are  seldom  found  in  the  pure  condition  indicated  by  the 
formulas,  but  are  often  mixed  with  sulphides  of  other 
substances  as  cobalt,  nickel,  silver  and  arsenic.  The  largest 
part  of  the  copper  of  commerce  comes  from  the  native 
ore  and  the  copper  pyrites.  The  most  productive  deposits 
are  those  of  Cornwall,  the  Siberian  mines  in  the  Ural 
Mountains,  the  mines  of  Chili,  of  southern  Australia,  of 
Saxony  and  the  Harz  region  in  Germany,  of  Spain  and 
especially  of  the  Lake  Superior,  the  Montana  and  the 
Arizona  mines  in  the  United  States.  At  the  present  time 
over  one  half  of  the  world's  supply  of  copper  is  produced 
in  this  country.  In  1896  the  total  production  was  420,000 
tons,  that  of  the  United  States  being  225,000  tons. 

History.  Copper  has  been  known  and  used  since  the 
earliest  historical  times  and  was  employed  largely  in  the 
production  of  bronze  by  the  Greeks,  Romans  and  Egyp- 
tians. Articles  of  bronze,  as  tools  and  weapons,  were  far 
more  common,  at  one  time,  than  those  of  iron  and  steel. 

Metallurgy.  The  production  of  pure  copper  from  the 
native  Lake  Superior  ore  is  largely  a  mechanical  opera- 
tion. The  ore  containing  the  free  copper  is  crushed  in 
stamp  mills  and  the  fine  product  so  obtained  is  systematic- 
ally washed  to  carry  away  the  lighter  rocky  or  earthy 
materials,  leaving  the  heavy  copper  behind.  Much  of  this 
native  copper  is  in  large  masses  which  simply  require 
fusion  to  be  brought  into  marketable  condition.  But  most 
of  the  metal  exists  in  a  finely  divided  condition  with  clay 
or  rock  and  this  requires  separation  by  crushing  and  wash- 
ing. The  thoroughly  washed  ore  is  finally  fused  with  a  little 
carbon  to  prevent  oxidation  and  cast  into  bars  or  ingots. 
This  copper  is  very  pure  and  is  valued  because  of  its 
freedom  from  other  metals. 

A  very  large  proportion  of  the  total  production  of  cop- 
per comes  from  the  several  sulphides  smelted.  The  reac- 
tions involved  are  somewhat  complicated,  but  may  be 
briefly  described  in  this  way:  The  ores  are  roasted  in  the 
air,  by  which  means  a  part  of  the  sulphur  is  expelled  as 


290  GENERAL  CHEMISTRY. 

SO2,  and  some  arsenic  and  other  impurities  are  lost.  This 
roasted  ore  is  then  fused  in  blast  furnaces  or  reverberatory 
furnaces  with  coke  and  slag  from  previous  operations. 
Much  of  the  iron  separates  as  silicate  in  the  slag  formed, 
while  the  copper  is  practically  all  left  as  sulphide.  There 
must  always  be  sulphur  enough  present  for  this  purpose, 
otherwise  copper  would  also  be  lost  in  the  slag.  This  latter 
does  not  mix  with  the  sulphide  readily  and  is  separated 
mechanically.  The  impure  copper  sulphide  is  called  the 
regulus  and  is  next  subjected  to  repeated  oxidations,  by 
which  most  of  the  sulphur  is  burned  away,  leaving  a  mix- 
ture of  Cu2O  and  Cu2S.  By  properly  conducting  the  oxi- 
dation the  proportions  of  these  can  be  controlled  so  that 
by  increasing  the  heat  of  the  furnace  they  will  react  on 
each  other  in  this  way: 

Cu2S-f2Cu2O=:6Cu+SO2. 

Metallic  copper  with  several  impurities,  and  known  as 
coarse  copper,  is  thus  secured.  This  is  melted  in  the  air  so 
that  the  traces  of  arsenic,  antimony,  sulphur,  iron,  lead  and 
other  substances  oxidize  and  escape  in  the  gaseous  condi- 
tion or  form  a  slag  with  the  silica  of  the  furnace  hearth, 
This  is  possible  because  all  these  elements  oxidize  very 
much  more  readily  than  does  copper.  However,  a  small 
amount  of  red  oxide  of  this  metal  is  formed  and  a  little 
sulphide  still  remains.  To  effect  final  purification  the 
molten  metal  is  covered  with  powdered  charcoal  and  is 
then  stirred  with  poles  of  green  wood,  which  gives  off 
water  vapor  and  reducing  gases.  These  combine  with  the 
oxygen  and  sulphur  and  leave  the  copper  in  practically 
pure  condition.  This  final  operation  is  known  as  poling 
the  copper. 

In  many  places  at  the  present  time  the  coarse  copper  is 
refined  by  a  process  of  electrolysis.  The  impure  metal  is 
cast  in  large  plates  which  are  hung  as  the  anodes  in  a  bath 
of  copper  sulphate  with  a  little  sulphuric  acid.  The  cath- 
odes consist  of  thin  sheets  of  copper.  On  passing  the 
current,  practically  pure  copper  deposits  on  these  sheets, 
which  thus  grow  up  into  bars.  The  copper  only  of  the 
anodes  is  taken  up,  as  lead  and  other  impurities  are  left 


GENERAL  CHEMISTRY.  291 

insoluble  or  oxidize  and  settle  to  the  bottom  of  the  re- 
fining vats.  This  electrolytic  copper  is  almost  absolutely 
chemically  pure. 

Properties.  Pure  copper  is  extremely  ductile  and 
malleable;  it  is,  next  to  silver,  the  best  conductor  of  heat 
and  electricity,  and  is  but  slightly  oxidized  in  either  dry 
or  moist  air.  It  alloys  readily  with  many  other  metals. 
Its  specific  gravity  is  over  8.9,  and  its  melting  point  is 
slightly  above  1,000°. 

Uses.  It  is  employed  largely  in  the  form  of  sheets  for 
making  water  baths,  air  baths,  stills  and  kettles  for  use  on 
the  large  and  small  scale.  Also  in  enormous  quantities  in 
the  form  of  wire  for  electric  conductors.  It  is  a  constit- 
uent of  several  valuable  alloys.  Bronze  is  an  alloy  of  cop- 
per and  tin  in  various  proportions  from  95  of  copper  and  5 
of  tin  to  75  of  copper  and  25  of  tin.  Some  kinds  of  bell 
metal  have  about  the  latter  composition.  Brass  is  an  alloy 
of  about  2  parts  of  copper  to  1  part  of  zinc.  Aluminum 
bronze  contains  usually  90  parts  of  copper  to  10  of  alumi- 
num. German  silver  is  an  alloy  of  copper,  zinc  and 
nickel. 

Copper  forms  two  classes  of  compounds,  known  as  cu- 
prous and  cupric.  In  the  first  the  metal  behaves  as  an 
univalent  element  while  in  the  second  the  condition  is  bi- 
valent. The  cupric  compounds  are  more  numerous  and  im- 
portant than  the  cuprous. 

Copper  Oxides.  Two  are  known,  the  cuprous  oxide 
or  red  oxide,  Cu2O,  and  the  black  cupric  oxide,  CtiO.  The 
first  is  prepared  by  heating  a  mixture  of  black  oxide  and 
copper  filings  in  equivalent  proportions,  or  by  the  reduc- 
tion of  an  alkaline  cupric  solution  by  means  of  a  solution 
of  dextrose,  which  was  illustrated  in  an  experiment  in  the 
chapter  on  carbon.  The  cuprous  oxide  is  bright  red  and 
is  employed  in  the  coloring  of  glass;  a  little  of  it  added 
to  the  ordinary  colorless  glass  in  the  melting  pots  con- 
verts it  into  ruby  glass.  The  black  oxide,  CuO,  may  be 
made  by  the  oxidation  of  metallic  copper  and  by  the  de- 


292  GENERAL  CHEMISTRY. 

composition  of  the  nitrate  by  heat.  It  is  sometimes  used 
in  coloring  glass  light  green,  and  is  employed  commonly 
as  an  oxidizing  agent  in  organic  combustion  analyses. 
This  oxide  dissolves  in  acids  yielding  the  cupric  salts, 
which  are  blue  or  green. 

Copper  Hydroxides.  Cuprous  hydroxide  is  known 
as  a  yellow  precipitate.  Cupric  hydroxide  is  obtained  as 
a  green  precipitate  by  adding  an  alkali  solution  to  a  solu- 
tion of  copper  sulphate  or  other  cupric  salt.  This  precipi- 
tate dissolves  in  ammonia  water  with  a  deep  blue  color.  It 
is  also  soluble  in  solutions  of  many  organic  substances  as 
the  sugars,  glycerol,  tartrates  and  mannitol.  The  solution 
in  Rochelle  salt  constitutes  the  ordinary  Fehling  solution 
employed  in  sugar  tests. 

Copper  Chlorides.  Cuprous  chloride,  CuCl,  or 
Cu2Cl2,  is  obtained  as  a  white  powder  insoluble  in  water 
by  heating  a  mixture  of  cupric  chloride  and  hydrochloric 
acid  with  scrap  copper;  on  dilution  jvith  water  the  white 
precipitate  appears.  It  dissolves  in  ammonia  and  this  so- 
lution has  important  applications  in  gas  analysis.  Cupric 
chloride,  CuCl2.2H2O,  is  a  green,  soluble  crystalline  com- 
pound, produced  by  dissolving  the  black  oxide  in  hydro- 
chloric acid.  It  is  also  used  in  analytical  chemistry. 

From  analogy  with  several  other  bodies  this  chloride  is 
often  considered  as  having  the  composition  CuCl,  but  a  de- 
termination of  the  vapor  density  fixes  the  molecular  weight 
as  198.2,  with  Cu2Cl2,  therefore,  as  the  formula. 

Copper  Sulphate.  The  most  important  of  the  copper 
compounds  is  blue  vitriol,  CuSO4.5H2O,  which  may  be 
made  by  the  solution  of  the  black  oxide  in  sulphuric  acid, 
but  is  also  largely  made  as  a  by-product  in  several  metal- 
lurgical processes.  It  is  readily  soluble  in  water  and  serves 
as  the  starting  point  in  the  preparation  of  several  other 
copper  compounds.  The  present  annual  production  of  the 
United  States  is  about  25,000  tons. 

A  solution  of  blue  vitriol  is  employed  in  copper  elec- 
troplating and  in  electrotyping,  the  metal  being  deposited 


•  GENERAL  CHEMISTRY.  293 

on  a  surface  by  electrolysis.  In  this  manner  a  firm  coating 
of  copper  may  be  put  on  iron  or  other  metals.  Or  a  coat- 
ing, which  may  be  stripped  off,  may  be  deposited  upon  a  wax 
or  plaster  of  Paris  cast,  thus  giving  a  copy  of  the  original. 
This  important  process  will  be  illustrated  by  an  experi- 
ment. 

Ex.  159.  Prepare  a  wax  copy  of  a  medal  in  the  following  manner: 
Melt  together  about  100  Gm.  of  beeswax,  5  Cc.  of  oil  of  turpentine  and 
10  Gm.  of  very  fine  graphite  powder.  Stir  thoroughly  to  distribute  the 
graphite  uniformly.  Next,  brush  fine  graphite  over  the  surface  of  the 
medal  to  be  copied,  and  then  wrap  a  strip  rrf  paper  around  its  edge  so  as 
to  form  a  cell  a  centimeter  or  more  in  depth,  the  bottom  of  which  is 
formed  by  the  medal;  support  the  paper  strip  with  a  rubber  band. 
Then  allow  the  wax  to  partially  cool,  stirring  it  meanwhile,  and  just 
before  it  begins  to  become  stiff  pour  enough  into  the  cell  to  half  fill  it. 
Then  lay  a  piece  of  copper  foil,  to  which  an  insulated  copper  wire 
about  15  Cm.  in  length  is  soldered,  on  the  wax  and  add  more  of  the 
melted  wax  to  quite  fill  the  cell,  which  is  next  set  aside  until  the  con- 
tents become  hard.  On  unwinding  the  strip  of  paper,  the  wax  may  be 
readily  detached  from  the  surface  of  the  medal,  and  on  its  under  surface 
will  be  found  a  sharp  impression  of  the  latter.  Sprinkle  some  fine  graph- 
ite on  this  wax  surface,  and  by  means  of  a  soft,  but  closely  filled  camel's 
hair  brush,  rub  the  graphite  over  the  whole  surface  so  as  to  impart  a 
high  polish  to  it.  If  the  wax  is  of  the  proper  degree  of  hardness  this 
can  be  easily  done  with  a  soft  brush,  having  short,  close  hair.  When 
the  polish  is  perfect  and  uniform  wash  off  the  loose  graphite  remaining, 
under  flowing  water,  and  pour  a  10  per  cent  copper  sulphate  solution 
over  the  polished  surface.  Distribute  this  evenly  by  means  of  a  very 
small  brush.  Then  add  some  fine  iron  filings,  free  from  grease,  and  mix 
them  thoroughly  with  the  copper  solution.  By  this  means  a  thin  coat- 
ing of  pure  copper  is  deposited  evenly  over  the  polished  surface  by  pre- 
cipitation. Add  more  solution  and  filings,  if  necessary,  to  make  this 
complete.  Then  wash  thoroughly  in  flowing  water  and,  without  delay, 
suspend  this  prepared  wax  copy  in  a  20  per  cent  solution  of  copper  sul- 
phate, slightly  acidulated  with  sulphuric  acid.  Attach  the  insulated 
copper  wire  leading  from  the  plate,  in  the  center  of  the  wax,  to  the  zinc 
pole  of  a  battery  of  five  or  six  Daniell  or  equivalent  elements  and  lead 
the  current  from  the  battery  into  the  solution  by  means  of  an  insulated 
wire  soldered  to  a  plate  of  copper.  Suspend  this  plate  so  that  it  stands 
parallel  to,  and  about  half  a  centimeter  away  from  the  coppered  surface 
of  the  wax  cast.  Under  these  conditions  copper  will  be  deposited  from 
the  solution  on  the  cast  and  taken  by  the  solution  from  the  copper  plate. 
At  the  end  of  about  24  hours  a  fine  layer  of  pure  copper  will  be  found 
on  the  wax,  and  can  be  readily  lifted  from  it.  It  will  be  seen  that  an 
accurate  copy  of  the  original  medal  has  been  made  by  the  copper,  the  wax 
copy  being  a  reversed  one.  Any  desired  thickness  of  the  copper  may  be 
deposited  by  giving  more  time  to  the  process.  The  addition  of  graphite 
to  the  wax  makes  it  a  conductor  of  electricity  and  prevents  its  perma- 


294  GENERAL  CHEMISTRY. 

nent  adherence  to  the  metal.  By  a  slight  modification  of  the  process 
illustrated,  an  electrotype  copy  of  a  page  of  type  may  be  made,  and  this 
may  be  used  in  place  of  the  real  type  in  printing. 

Blue  vitriol  when  carefully  heated  loses  all  of  its 
water  of  crystallization  and  leaves  a  white  powder  of  the 
pure  sulphate,  CuSO4.  When  this  is  exposed  to  moist  air 
it  absorbs  water  and  becomes  blue.  It  is  also  employed 
as  a  test  substance  for  the  presence  of  water  in  alcohol 
and  certain  other  organic  liquids.  In  absolute  alcohol  the 
anhydrous  sulphate  remains  white,  but  if  a  very  small 
amount,  even,  of  water  is  present,  it  becomes  blue  after  a 
time.  Ammonia  is  absorbed  by  anhydrous  copper  sulphate, 
forming  a  blue  compound  with  the  formula  CuSO4.5NH3, 
which  by  heat  is  changed  to  CuSO4.NH3.  A  crystalline 
compound,  CuSO4.4NH3.H2O,  is  also  known. 

Copper  Arsenite  is  formed  by  precipitating  a  solution 
of  copper  sulphate  with  one  of  sodium  or  potassium 
arsenite.  This  is  a  very  poisonous  substance  and  is  com- 
monly known  as  Scheele's  green.  A  somewhat  similar 
green  precipitate  is  formed  by  mixing  a  boiling  solution  of 
copper  acetate  with  a  boiling  solution  of  arsenous  acid. 
This  product  is  known  as  Schweinfurt  green. 

Recognition.  Copper  compounds  are  easily  recog- 
nized by  the  precipitation  of  the  metal  on  a  strip  of  bright 
iron  from  a  solution  of  the  sulphate  or  other  salt,  by  the 
deep  blue  color  produced  by  the  addition  of  ammonia  in 
excess  to  copper  solutions,  and  finally  by  the  red  precipitate 
of  copper  ferrocyanide,  Cu2FeC6N6,  produced  on  the 
addition  of  potassium  ferrocyanide  to  a  neutral  or  slightly 
acid  copper  solution. 

SILVER. 

Occurrence.  Silver  is  found  in  small  quantities  in  the 
uncombined  condition,  but  it  usually  occurs  as  sulphide 
with  lead,  arsenic  and  antimony,  and  occasionally  as 
chloride.  Of  the  sulphides  silver  glance  ore,  Ag2S,  red 
silver  ore,  Ag3AsS3,  and  brittle  silver  ore,  5Ag2S.Sb2S3, 


GENERAL  CHEMISTRY.  295 

are  the  most  important;  but  many  deposits  of  galena, 
PbS,  contain  the  silver  as  sulphide,  and  much  of  the  silver 
of  commerce  comes  from  this  source. 

The  most  productive  silver  regions  of  the  world  are 
those  of  the  western  United  States,  Mexico,  Bolivia  and 
Australia. 

The  world's  total  production  of  silver  in  1896  was 
5,786,567  kilograms.  Of  this  the  United  States  furnished 
1,819,208  kilograms,  Mexico  1,286,842  kilograms,  Bolivia 
638,000  kilograms,  Australia  605,400  kilograms  and 
Germany  428,429  kilograms. 

History.  Silver  has  been  known  from  the  earliest 
times,  and  has  always  been  highly  prized  on  account  of  its 
color  and  valuable  property  of  alloying  with  other  metals. 
The  amount  of  silver  known  to  the  ancients  was,  however, 
very  small  as  compared  with  the  amounts  now  in  use. 

Metallurgy.  Silver  is  separated  from  its  ores  by  sev- 
eral methods,  of  which  the  two  most  important  will  be 
briefly  described. 

Amalgamation  Process.  This  is  a  process  in  which  sil- 
ver is  taken  up  by  mercury  from  the  ore  after  preliminary 
treatment.  This  treatment  varies  in  different  localities, 
and  in  many  cases  no  simple  explanation  can  be  given  of 
the  chemical  principles  involved  in  it.  Often  the  ore  is 
crushed  and  roasted  with  common  salt  to  convert  sulphide 
into  chloride  of  silver.  The  finely  ground  product  is  then 
agitated  with  water  and  scrap  iron,  by  which  the  chloride 
of  silver  is  reduced  to  the  metallic  condition.  Mercury  is 
added  and  thoroughly  mixed  with  the  mud  for  several 
hours,  until  the  amalgamation  or  solution  of  the  silver  in 
the  mercury  is  practically  complete.  This  operation  is 
carried  out  in  large  tanks,  and  when  finished  water  is  run 
in  to  wash  away  the  lighter  mud,  leaving  the  heavy  amal- 
gam at  the  bottom.  The  amalgam  is  gathered  up,  pressed 
into  balls,  dried  and  distilled  from  a  retort,  by  which 
means  the  mercury  is  recovered  for  further  use  and  the 
silver  left  in  metallic  condition,  ready  for  the  refiner.  In 
some  cases  the  amalgamation  is  conducted  without  the 


296  GENERAL  CHEMISTRY. 

preliminary  roasting  of  the  ore,  and  it  is  also  possible  to 
reduce  the  silver  chloride  in  the  amalgamation  process  it- 
self, without  the  use  of  scrap  iron.  As  silver  ores  often 
contain  gold,  this  metal  will  be  left  with  the  crude  silver 
in  the  retort.  The  method  of  refining  will  be  explained 
under  gold. 

Smelting  Process.  As  silver  ores  usually  contain  lead  it 
has  been  found  most  economical  and  simple  to  combine  the 
reduction  of  the  two  metals  by  a  smelting  process,  the  im- 
portant principle  involved  being  this:  Lead  ores,  by  the 
addition  of  proper  fluxes  may  be  easily  reduced  in  a  blast 
furnace  by  coke  to  the  metallic  condition.  This  reduced 
lead  in  turn  acts  on  the  silver  and  gold  compounds,  reduc- 
ing them  likewise  to  metals  and  alloying  with  them.  The 
molten  lead  then  settling  to  the  bottom  of  the  blast  fur- 
nace carries  the  other  metals,  and  from  time  to  time  is  run 
off  into  molds,  forming  pigs,  from  which  the  metals  pres- 
ent may  be  separated  by  subsequent  operations.  The 
smelting  operation  is  usually  preceded  by  a  roasting  to  re- 
move a  large  part  of  the  sulphur  and  form  lead  oxide.  It 
is  customary  to  charge  the  blast  furnace  with  a  mixture  of 
ores  rather  than  with  a  single  one,  in  order  to  secure  by 
the  proper  combination  of  acid  and  basic  material  a  charge 
which  will  flux  easily,  that  is,  in  which  the  heavy  metals 
will  separate  perfectly  from  the  other  substances  left  as  a 
slag. 

The  crude  lead  from  the  smelting  furnace  is  known  as 
base  bullion  and  is  treated  for  separation  of  the  several 
metals  in  it.  The  principal  steps  in  the  process  can  be 
only  briefly  described.  The  bullion  is  heated  on  the  hearth 
of  a  reverberatory  furnace  and  stirred  thoroughly  to  bring 
every  part  in  contact  with  air.  This  suffices  to  oxidize 
several  metals  which  may  be  present  in  small  amount,  and 
separate  them  as  a  dross  which  may  be  removed  by  skim- 
ming. Antimony,  copper  and  arsenic  are  so  separated. 
What  is  left  is  generally  treated  by  the  Parkes  process. 
The  metal  is  melted  in  large  iron  pots  and  a  small  amount 
of  zinc,  1  to  2  per  cent  by  weight,  is  added  and  thor- 
oughly incorporated  by  stirring.  Zinc  combines  readily 
with  silver  and  gold,  but  only  in  limited  proportion  with 


GENERAL  CHEMISTRY.  29? 

lead.  The  zinc  therefore  separates,  comes  to  the  surface, 
carrying  with  it  the  precious  metals.  The  whole  mass  is 
allowed  to  cool  slowly  and  the  zinc  alloy  on  the  surface  be- 
gins to  solidify  first.  It  is  skimmed  off  with  a  perforated 
ladle  and  transferred  to  a  separate  vessel.  To  the  residue 
of  lead  more  zinc  may  be  added,  a  second  and  third  time 
if  necessary,  which  is  determined  by  the  richness  in  silver 
of  the  original  bullion,  the  main  portion  of  which  is  lead, 
of  course.  A  very  little  of  the  zinc  remains  with  the  lead, 
while  the  other  metals  are  practically  all  removed.  From 
the  large  mass  of  lead  the  trace  of  zinc  is  removed  by 
heating  to  a  high  temperature,  as  zinc  volatilizes  readily.  It 
remains  to  separate  the  precious  metals.  To  this  end  the 
zinc  alloy  is  heated  in  graphite  crucibles  or  retorts  to  such 
a  temperature  that  the  zinc  distills  out  almost  completely 
and  is  condensed  for  use  again.  The  residue  contains  lead 
with  the  precious  metals  and  the  former  is  eliminated  by 
oxidation  in  a  current  of  hot  air,  the  lead  being  converted 
into  litharge,  a  little  of  which  is  absorbed  by  the  hearth 
of  the  cupel  furnace  in  which  the  oxidation  is  conducted, 
the  most  being  floated  off  as  a  molten  slag  from  the  top  of 
the  heavier  metal.  This  litharge  is  later  reduced  to  lead. 

The  metal  now  remaining  consists  of  silver  mainly, 
with  a  little  gold,  and  possibly  traces  of  other  metals.  The 
separation  of  these  will  be  described  a  few  pages  in  ad- 
vance, under  gold.  It  has  been  recently  discovered  that  a 
minute  amount  of  aluminum  added  with  the  zinc  in  this 
process  aids  materially  in  the  separation  of  the  silver  and 
gold  from  the  lead.  Other  methods  and  modifications  are 
used  in  various  places;  but  the  details  cannot  be  explained 
here. 

Properties.  Silver  is  the  best  known  conductor  of 
heat  and  electricity,  is  very  malleable  and  ductile,  and  al- 
loys readily  with  many  metals.  At  a  high  temperature  it 
absorbs  gases  in  large  volume,  to  give  them  out  on  cool- 
ing. The  metal  is  not  dissolved  by  hydrochloric  or  dilute 
sulphuric  acid,  but  is  dissolved  by  strong  hot  sulphuric 
acid  and  by  nitric  acid. 


298  GENERAL  CHEMISTRY. 

Uses.  The  metal  is  largely  used  in  coinage,  alloyed 
with  copper  to  harden  it.  In  this  country  the  alloy  con- 
tains 90  parts  of  silver  to  10  of  copper.  In  solid  silverware 
the  metal  is  alloyed  in  the  same  manner,  but  often  with 
more  copper.  Immense  quantities  of  silver  are  employed 
in  plating  by  a  process  essentially  similar  to  that  illustrated 
under  copper.  The  silver  plating  bath  is  a  solution  of 
silver  cyanide  in  an  excess  of  potassium  cyanide,  and  is 
made  in  various  ways.  Any  amount  of  silver  may  be  de- 
posited on  copper,  brass,  pewter  and  certain  other  alloys, 
but  in  practice  the  thickness  of  the  film  is  not  great.  A 
film  having  a  thickness  of  0.01  Mm.  is  practically  a  good 
one.  This  contains  about  1  gram  of  silver  to  the  square 
decimeter.  Three  grams  of  silver  on  a  square  decimeter 
is  a  very  good  plate  and  has  the  thickness  of  thin  writing 
paper. 

All  articles  to  be  covered  must  be  thoroughly  cleaned 
first,  by  immersion  in  a  hot  alkali  solution  in  case  of  Bri- 
tannia metal  or  pewter,  and  in  nitric  acid  in  the  case  of 
copper  or  brass.  These  baths  remove  the  impurities  on 
the  surfaces  to  be  covered.  The  articles  are  then  thor- 
oughly washed  and  suspended  in  the  plating  solution.  But 
many  details  in  the  process  of  washing  which  cannot  be 
explained  here  must  be  observed.  It  is  absolutely  neces- 
sary to  have  the  surface  perfectly  clean  and  free  from 
oxide,  scale  or  impurity  of  any  kind. 

Plating  Solution. 

As  it  may  interest  the  student  to  have  the  composition 
of  a  good  plating  bath,  the  following  figures  are  given  : 
Dissolve  17  Gm.  of  silver  nitrate  in  about  150  Cc.  of  water. 
Make  a  solution  of  10  Gm.  of  commercial  potassium  cya- 
nide in  100  Cc.of  water  and  mix  the  two  solutions,  stirring 
well.  Allow  to  settle  and  then  pour  off  two  portions  of  the 
clear  liquid  of  about  25  Cc.  each.  To  one  add  a  few  drops 
of  dilute  silver  nitrate  solution.  If  a  precipitate  forms  here 
it  shows  that  too  much  cyanide  had  been  added.  In  this 
case  return  both  portions  to  the  main  liquid  and  add  a 
little  more  silver  solution  as  long  as  a  precipitate  appears 


GENERAL  CHEMISTRY.  299 

to  form.  On  the  other  hand,  if  the  silver  nitrate  solution 
does  not  make  a  precipitate,  add  a  little  cyanide  solution 
to  the  second  small  portion  of  liquid  poured.  A  precipitate 
should  show  here,  indicating  an  excess  of  silver  in  the 
mixture.  Return  both  small  portions  to  the  original,  and 
after  settling  add  gradually,  and  a  very  little  at  a  time, 
some  cyanide  solution  as  long  as  a  precipitate  forms.  Be 
careful  not  to  add  more  than  this.  Allow  the  precipitate 
to  settle  thoroughly,  pour  off  the  liquid,  and  wash  the  resi- 
due by  decantation  several  times.  This  gives  us  silver 
cyanide  in  nearly  pure  condition.  Now  dissolve  15  to  20 
Gm.  of  commercial  potassium  cyanide  in  200  Cc.  of  water 
and  add  this  to  the  moist  silver  cyanide  in  a  beaker  or 
bottle.  The  latter  goes  into  solution,  forming  the  desired 
double  cyanide  with  an  excess  of  potassium  cyanide  left. 
Finally,  dilute  the  whole  with  water  to  make  a  liter. 

The  articles  to  be  plated  are  hung  from  copper  wires 
into  this  bath  and  attached  to  the  zinc  pole  of  a  battery  of 
Bunsen  or  chromate  cells.  The  current  is  led  into  the 
solution  from  the  other  pole  through  a  plate  of  metallic 
silver.  By  this  arrangement  silver  dissolves  from  the 
plate  as  fast  as  it  is  deposited  on  the  article,  thus  keeping 
the  solution  of  nearly  constant  strength.  To  insure  an 
even  deposit  the  article  must  be  moved  frequently  to  bring 
all  sides,  in  turn,  near  the  silver  electrode. 

The  deposit  thus  formed  has  a  rough  or  "matt"  sur- 
face and  can  be  given  the  usual  smooth  finish  by  rubbing 
with  chalk  or  other  polishing  powder.  On  removing  the 
plated  article  from  the  bath  it  is  washed  thoroughly  in 
clean  water  and  dried  in  an  atmosphere  free  from  tarnish- 
ing gases. 

Compounds  of  Silver.  Numerous  compounds  of  this 
metal  are  known,  but  the  most  important  are  the  nitrate, 
AgNO3,  the  chloride,  AgCl,  the  bromide,  AgBr,  and  the 
iodide,  Agl. 

Silver  Nitrate.  This  salt  is  made  in  large  quantities 
by  dissolving  the  metal  in  nitric  acid.  It  is  very  soluble 
in  water  and  is  obtained  crystallized  in  plates.  The  im- 


300  GENERAL  CHEMISTRY. 

portant  uses  of  the  nitrate  are  in  the  preparation  of  plating 
solutions,  as  just  described,  and  in  making  certain  salts  em- 
ployed in  photography;  under  the  name  of  lunar  caustic  it 
is  also  used  in  medicine.  It  is  readily  obtained  in  a  state 
of  great  purity. 

Silver  Chloride,  Bromide  and  Iodide.  These  three 
substances  are  interesting  mainly  because  of  their  behav- 
ior in  light.  On  this  behavior  the  ordinary  processes  of 
photography  are  founded.  Some  experiments  will  be  given 
illustrating  these  points. 

Ex.  160.  In  each  of  three  test-tubes  take  about  5  Cc.  of  a 
dilute  solution  of  silver  nitrate.  Add  to  one  a  few  Cc.  of  a  solution  of 
sodium  chloride,  to  the  second  a  solution  of  sodium  bromide  and  to  the 
third  some  dilute  potassium  iodide  solution.  In  each  case  a  curdy  pre- 
cipitate forms  which  is  most  characteristic  after  shaking.  These  pre- 
cipitates are  the  chloride,  the  bromide  and  the  iodide  of  silver,  of  which 
the  first  is  white  and  the  others  yellowish  white  in  dilute  mixture. 
Divide  the  contents  of  each  test-tube  into  three  parts.  To  one  part,  in 
each  case,  add  ammonia  water.  The  precipitates  of  silver  chloride 
and  bromide  dissolve  while  the  iodide  is  found  to  be  insoluble. 
It  will  be  noticed  that  the  chloride  is  much  more  easily  soluble 
than  the  bromide.  To  another  portion,  from  each  one  of  the  precipi- 
tated mixtures,  add  some  solution  of  sodium  thiosulphate.  The  three 
precipitates  dissolve.  Next,  stand  the  remaining  portions  in  bright  sun- 
light. After  a  time  it  will  be  noticed  that  they  darken  and  become  vio- 
let in  shade.  The  extent  of  this  darkening  can  be  materially  increased 
by  having  a  slight  excess  of  silver  nitrate  present  in  each  case,  that  is,  a 
little  more  than  can  be  precipitated  by  the  salts  added.  After  the  three 
tubes  have  stood  in  the  light  long  enough  for  their  contents  to  darken 
throughout,  which  is  aided  by  shaking  occasionally,  repeat  the  experi- 
ment of  adding  sodium  thiosulphate.  It  will  now  be  found  that  the  pre- 
cipitates have  become  insoluble.  The  action  of  the  light  has  converted 
them  into  compounds  which  can  no  longer  dissolve  in  a  liquid  which  was 
a  good  solvent  for  the  fresh  precipitates. 

The  precipitations  follow  according  to  these  equations: 
AgNO3+NaCl  =AgCl  +NaNO3 


AgN03  +  KI      =AgI    +KN03. 

The  experiment  shows  that  in  fresh  condition  these  pre- 
cipitates are  soluble  in  a  certain  solution,  that  they  are 
changed  in  some  manner  by  exposure  to  light,  and  that 
after  this  change  has  taken  place  they  are  no  longer  solu- 


GENERAL  CHEMISTRY.  301 

ble  in  the  solution  used.  How  this  fact  can  be  practically 
applied  will  be  shown  below.  We  need  for  the  purpose 
what  is  called  sensitive  paper,  that  is,  paper  which  holds 
on  one  surface  a  precipitate  as  above  described.  Such  a 
paper  may  be  prepared  as  follows: 

Silver  Paper. 

Dissolve  a  gram  of  ammonium  chloride  in  15  Cc.  of 
water,  and  add  about  1  Cc.  of  alcohol.  To  this  solution 
add  50  Cc.  of  white  of  egg.  The  white  of  egg  for  the 
purpose  is  best  made  by  separating  the  yolk  mechanically, 
and  shaking  the  white  portion  thoroughly  with  broken  glass 
in  a  bottle.  The  mixture  is  allowed  to  settle,  and  then  the 
liquid  portion  is  filtered  through  well  washed  cotton  into 
the  other  solution,  as  mentioned.  The  mixture  is  shaken 
and  poured  out  into  a  shallow  dish.  On  this  liquid  sheets 
of  pure  white  slightly  glazed  paper,  free  from  chemical 
bleaching  agents,  are  floated  a  few  minutes,  and  then  hung 
up  to  drain  and  dry.  The  paper  so  prepared  is  next  floated 
on  a'solution  containing  10  Gm.  of  silver  nitrate  in  100  Cc. 
of  water.  This  must  be  done  in  a  darkened  room.  The 
paper  is  left  on  the  solution  three  to  five  minutes.  A  pre- 
cipitate of  silver  chloride  in  presence  of  an  excess  of  silver 
nitrate  forms.  The  paper  so  impregnated  is  hung  up 
again  in  a  dark  room  to  drain  and  become  dry.  It  is  then 
kept  in  a  box  away  from  the  light  until  used.  The  student 
may  employ  such  paper,  but  it  will  generally  be  found 
more  convenient  to  purchase  similar  paper  from  dealers  in 
photographic  supplies.  With  such  paper  the  following 
simple  experiment  may  be  made  : 

Photography.     Silver  Printing. 

Ex.  161.  Cut  out  a  small  figure  of  any  shape  from  a  piece  of  dark 
paper  and  spread  this  over  the  center  of  a  slightly  larger  piece  of  the 
above  sensitive  paper,  and  on  the  side  holding  the  precipitate.  Then 
place  these  between  two  pieces  of  clean,  colorless  glass,  which  may  be 
clamped  together  at  opposite  corners  by  means  of  bits  of  brass  wire. 
Expose  the  paper  so  held  to  the  action  of  sunlight,  the  sensitive  surface 
facing  the  light.  The  glass  over  the  other  surface  should  be  covered  by 


302  GENERAL  CHEMISTRY. 

some  opaque  object  to  prevent  the  light  from  reaching  the  paper  from 
this  direction.  After  an  exposure  of  ten  minutes  remove  the  glass  in  a 
part  of  the  room  away  from  bright  light  and  observe  that  a  white  image 
of  the  object,  represented  by  the  cut  opaque  paper,  is  found  on  a  dark 
ground.  Wash  the  paper  a  minute  in  clean  water  and  then  immerse  it 
for  twenty  minutes,  with  frequent  stirring,  in  a  solution  containing  10 
Gm.  of  sodium  thiosulphate  in  100  Cc.  of  water.  After  this  treatment 
transfer  the  paper  to  clean  water  and  allow  it  to  soak  an  hour,  with  fre- 
quent changes  of  the  water.  Then  hang  it  up  or  put  it  between  blotting 
paper  to  dry.  We  have  now  a  permanent  silver  print  which  may  be 
exposed  to  the  light  without  further  change. 

Next,  repeat  the  above  experiment,  but  do  not  wash  the  print  in 
water  or  the  sodium  thiosulphate  solution.  It  will  be  found  that  the 
print  thus  made  is  not  permanent,  even  in  a  faint  light.  After  a  few 
days  the  light  part  will  grow  dark  and  in  time  the  whole  surface  will  be 
uniform  in  color,  with  complete  loss  of  the  image.  Pieces  of  paper  5  Cm. 
square  are  large  enough  for  the  purpose. 

The  operations  described  above  represent  those  carried 
out  in  printing  and  fixing  (making  permanent)  a  photo- 
graph on  paper.  In  practical  work  the  sensitive  paper  is 
exposed  behind  a  glass  "  negative,"  that  is,  a  picture  in 
light  and  shade,  on  glass,  made  by  a  process  which  cannot 
be  well  illustrated  in  all  details  in  the  ordinary  laboratory. 
It  may  be  said,  however,  of  this  process  that  the  main 
chemical  principles  involved  are  the  same  as  in  printing 
and  fixing  the  photograph. 

Glass  Negatives. 

As  employed  in  ordinary  dry- plate  photography,  the 
negative  plate  is  usually  a  sheet  of  glass  on  which  the  sen- 
sitive silver  compound  is  held  in  the  condition  of  a  fine 
emulsion.  Such  an  emulsion  may  be  prepared  by  soaking 
gelatin  in  water  to  soften  it,  and  then  heating  the  mixture 
until  homogeneous.  To  this,  potassium  bromide  in  solu- 
tion is  added  and  thoroughly  distributed  by  stirring. 
Then  a  certain  weight  of  silver  nitrate  in  solution  is  poured 
in  and  thoroughly  stirred  to  give  a  precipitate  of  silver 
bromide,  which  is  very  fine  and  evenly  distributed  through 
the  whole  mass.  This  is  allowed  to  set  and  is  then  cut 
into  thin  strips,  which  are  soaked  in  water  until  all  excess 
of  soluble  bromide  is  washed  out.  What  is  left  is  softened 
in  pure  water,  by  aid  of  heat,  and  in  this  condition  poured 


GENERAL  CHEMISTRY.  303 

over  clean  glass  plates,  which  are  then  allowed  to  stand  on 
a  level  shelf  in  a  place  free  from  dust  until  the  gelatin  sets 
again.  All  of  these  operations  must  be  performed  in  a 
nearly  dark  room,  illuminated  by  faint  nonactinic  light 
only.  The  plates  are  afterward  dried  by  warm  air,  after 
most  of  the  moisture  has  evaporated  at  a  low  temperature. 

Such  plates  when  properly  made  are  very  sensitive, 
and  may  be  kept  indefinitely  in  the  dark.  When  exposed 
in  the  camera  the  sensitive  silver  bromide  is  acted  on  by 
the  light  reflected  from  white  or  colored  objects.  The 
extent  of  the  decomposition  of  the  bromide  is  proportional 
to  the  intensity  of  the  light  reaching  the  plate  through  the 
lens  of  the  camera;  but  when  the  plate  is  taken  from  its 
holder,  in  the  dark  room,  and  examined  by  a  faint  red 
light  no  change  can  be  seen  on  it.  It  contains  the  ele- 
ments of  an  image,  but  this  is  in  a  latent  condition,  and 
must  be  brought  out  by  the  aid  of  what  is  called  a  developer. 
This  is  a  solution  of  some  reducing  or  oxygen  absorbing 
substance,  such  as  hydroquinon,  ferrous  sulphate,  pyro- 
gallol  (called  pyrogallic  acid)  and  other  bodies,  dissolved 
with  alkalies  and  certain  additions  which  need  not  be  men- 
tioned here.  On  immersing  the  plate  in  the  developing 
solution  an  image  soon  becomes  visible,  and  this  is  formed 
by  the  precipitation  of  metallic  silver  as  a  film  of  greater  or 
less  thickness,  producing  variations  in  light  and  shade 
which  may  be  compared  to  the  heavy  and  light  crayon 
strokes,  properly  applied,  that  are  sufficient  to  make  an 
image  on  a  piece  of  white  paper. 

Wherever  the  light  has  acted  strongly  on  the  plate  the 
development  produces  a  relatively  heavy  precipitate  of 
silver.  Where  the  light  has  been  thrown  with  less  inten- 
sity the  silver  deposit  has  less  depth.  We  are  not  able  to 
explain  the  exact  nature  of  the  change  which  the  light 
brings  about  in  the  silver  bromide  on  the  plate,  but  it  is 
certainly  left  in  a  condition  in  which  reduction  to  the  state 
of  metal  is  relatively  rapid.  The  reduction  of  the  ordinary 
bromide,  AgBr,  is  much  slower.  The  action  of  the  hydro- 
quinon developer  is  illustrated  by  this  equation  : 

2AgBr+C6H402H2+2KOH  = 

2Ag+2KBr-f2H2O-fC6H4Og. 


304  GENERAL  CHEMISTRY. 

Hydroquinon,  C6H4O2H2,  in  reducing  the  silver  salt 
becomes  oxidized  to  quinon,  CCH4O2.  Addition  of  alkali 
to  the  developer  hastens  the  action,  and  addition  of  bro- 
mide retards  it,  as  it  is  a  product  of  the  reaction. 

The  photographer  allows  the  plate  to  remain  in  the 
developing  solution  long  enough  to  produce  a  good  image, 
which  is  determined  by  frequent  examination  by  aid  of  a 
faint  red  light.  The  plate  is  then  washed  in  water  a  short 
time,  and  immersed  next  in  the  solution  of  sodium  thiosul- 
phate,  as  already  described.  This  removes  the  silver  salt 
present  which  is  in  excess  of  the  amount  needed  to  make 
the  picture.  Finally,  the  negative  is  thoroughly  washed  in 
water,  and  then  stood  on  its  edge  to  drain  and  become  dry. 
It  is  after  this  ready  for  use  to  make  "prints"  or  paper 
copies,  as  already  illustrated.  The  sensitive  paper  and  the 
glass  negative  are  placed  face  to  face  in  a  printing  frame, 
which  is  then  exposed  so  that  light  may  shine  through  the 
plain  glass  surface  toward  the  paper.  The  dark  parts  of 
the  negative,  however,  will  not  allow  light  to  pass,  and 
consequently  the  paper  remains  unacted  upon  at  such 
points.  On  the  other  hand,  it  is  strongly  attacked  where 
the  silver  deposit  on  the  negative  is  thin,  and  an  image  is 
thus  made  on  the  paper  corresponding  to  that  on  the  glass, 
but  with  the  lights  and  shades  reversed.  Light  objects, 
therefore,  which  appear  dark  in  the  "negative"  are  prop- 
erly represented  in  the  print  or  "positive"  picture. 

The  action  of  the  sodium  thiosulphate  as  a  fixing  agent 
depends  on  the  fact  that  it  dissolves  silver  bromide,  or 
chloride,  forming  a  soluble  double  salt. 


Na2S2O3+AgBr=:AgNaS2O3 

Only  the  excess  of  bromide  is  removed  in  this  manner. 
By  sufficient  washing  the  AgNaS2O3  may  be  wholly  elimi- 
nated. 

Among  other  silver  compounds  there  may  be  men- 
tioned the  oxide,  Ag2O,  which  with  water  behaves  as  an 
alkali,  AgOH;  the  phosphate,  Ag3PO4,  a  yellow  precipi- 
tate; the  chromate,  Ag2CrO4,  a  red  precipitate,  and  the 
thiocyanate,  AgSCN,  a  white  precipitate,  all  of  interest  in 
analytical  chemistry. 


GENERAL  CHEMISTRY.  305 

Recognition.  Silver  compounds  are  recognized  by  the 
formation  of  the  precipitates  just  mentioned,  and  com- 
monly by  the  precipitation  of  silver  chloride  on  addition  of 
hydrochloric  acid  to  a  silver  solution.  This  precipitate 
dissolves  in  ammonia,  in  potassium  cyanide  solution,  and 
in  a  thiosulphate  solution. 

GOLD. 

Occurrence.  The  metal  is  usually  found  in  the  native 
condition  disseminated  through  sand  or  quartz  rocks  or 
with  pyrites.  It  may  occur  in  the  old  eruptive  rocks  or 
veins,  or  in  the  disintegration  products  of  such  rocks,  that 
is  in  sands  or  conglomerate  masses.  The  chief  gold  pro- 
ducing countries  are  the  United  States,  Dutch  and  British 
South  Africa,  Australia,  Russia  and  British  North  America. 
In  1896  the  world's  production  of  gold  was  316,254  kilo- 
grams, of  which  this  country  furnished  79,576  kilograms, 
or  over  one-fourth. 

History.  We  find  gold  mentioned  in  the  earliest 
historical  accounts,  and  it  has  always  ranked  as  a  standard 
of  value.  The  gold  known  to  the  ancients  was  usually 
obtained  by  the  washing  of  gold  bearing  sands  from  the 
beds  of  rivers.  The  methods  of  separating  it  from  rock 
masses  are  comparatively  recent. 

Metallurgy.  Much  that  was  said  about  silver  applies 
to  gold.  When  gold  bearing  sands  are  shaken  in  an  iron 
pan  under  a  stream  of  water  the  lighter  earthy  particles 
are  washed  away  and  the  heavy  gold  remains.  In  a  modi- 
fication of  this  process  great  banks  of  sand  are  washed 
down  into  long  sluice  boxes  lined  with  amalgamated  copper 
sheets.  The  heavy  gold  lags  behind,  much  of  it  being 
held  by  the  mercury,  while  the  lighter  sands  are  gradually 
carried  away  by  the  running  water  in  the  sluices,  which 
may  be  hundreds  of  feet  in  length.  The  gold  is  scraped 
up  from  the  bottom  of  the  boxes  and  refined.  Where  gold 
occurs  in  masses  of  solid  rock  such  processes  will  not 
suffice  and  it  is  necessary  to  resort  to  others. 

Amalgamation  Process.     In  this  the  rock  is  crushed  and 


306  GENERAL  CHEMISTRY. 

ground  to  a  fine  powder  in  stamp  mills.  It  is  washed  from 
under  the  stamps  by  a  stream  of  running  water  into  large 
vats  where  it  is  agitated  with  mercury,  which  gathers  up 
the  gold  particles  from  the  mud  and  dissolves  them  as  in 
the  case  of  silver.  This  amalgam  is  afterwards  collected 
and  distilled,  the  mercury  being  saved  and  the  gold  left  in 
impure  condition.  It  is  melted  into  bars  and  sent  to  the 
refiner. 

Smelting  Process.  Many  ores  cannot  be  economically 
or  completely  extracted  by  the  mercury  method  and  these 
are  often  treated  by  smelting  with  lead  ores,  exactly  as  de- 
scribed for  silver.  The  treatment  of  the  base  bullion  ob- 
tained is  the  same  as  there  given.  At  the  present  time  a 
large  part  of  the  rich  ore  in  this  country  is  so  handled. 

Cyanide  Process.  Metallic  gold  is  dissolved  by  solutions 
of  potassium  cyanide  under  certain  condi'tions,  and  a  pro- 
cess suggested  years  ago  has  recently  been  revived  by 
which  many  of  the  poorer  gold  ores  are  successfully 
treated  in  great  quantities.  The  ores  are  crushed,  some- 
times after  preliminary  roasting,  and  percolated  in  tanks 
with  a  weak  solution  of  potassium  cyanide  to  which  sodium 
dioxide  is  sometimes  added  to  hasten  oxidation.  The  re- 
actions taking  place  are  somewhat  complex,  but  the  most 
probable  ones  are  these  : 


=  2AuK(CN)8+2KOH 
Au+2KCN-f-H2O        =  AuK(CN)2-f  KOH  +  H. 

Silver  in  the  ores  dissolves  in  the  same  manner.  By 
this  solvent  action  it  is  possible  to  dissolve  traces  of  the 
precious  metals  from  large  quantities  of  sand  or  crushed 
rock.  From  the  solutions  obtained  the  gold  and  silver  are 
precipitated  by  electricity  or  by  addition  of  metallic  zinc: 

2KAu(CN)2-fZn=:K2Zn(CN)4  +  2Au. 

The  gold  settles  to  the  bottom  of  the  precipitation 
tanks  as  a  fine  slime.  It  is  collected,  dried,  smelted  and 
refined. 

Chlorination  and  Bromination  Processes.       In  these  the 
ores  are  roasted  to  destioy  sulphides  and  then  subjected  in 


GENERAL  CHEMISTRY.  307 

casks  to  the  action  of  chlorine  gas  or  bromine  vapor.  The 
gold  dissolves  to  form  the  chloride,  AuCl3,  or  bromide, 
which  is  readily  soluble  in  water.  At  the  end  of  the  treat- 
ment water  is  allowed  to  percolate  through  the  casks  to 
wash  out  the  soluble  salt.  From  this  solution  the  gold  is 
precipitated  by  charcoal  powder  or  by  ferrous  sulphate  : 

AuCl3  +  3FeSO4=Au+Feg(SO4)3  +  FeCl3. 

The  gold  settles  out  as  a  fine,  dark  powder,  which  is 
collected,  dried  and  refined. 

Refining  of  Gold.  The  crude  gold  bullion  obtained 
by  any  one  of  the  above  processes  contains  silver  usually, 
and  possibly  small  amounts  of  copper  and  other  metals  ; 
it  must  therefore  be  subjected  to  a  separating  or  refining 
process.  For  many  years  the  following  method  was  com- 
monly followed,  and  is  yet  to  some  extent.  The  metal  is 
assayed  and,  if  rich  in  gold,  enough  silver  is  melted  in  with 
it  to  make  the  gold  about  one-fourth  of  the  whole.  Such 
an  alloy  is  readily  attacked  by  nitric  acid,  while  in  alloys 
much  richer  in  gold  the  action  is  very  slow  or  may  not  take 
place  at  all.  In  the  weak  alloy  all  the  baser  metals  are 
easily  dissolved  as  nitrates,  while  the  gold  is  left  as  a  dark, 
spongy  mass,  which  maybe  washed  quite  pure  with  water. 
After  this  treatment  it  is  dried  and  melted  into  bars  of 
fine  gold.  The  mixed  nitrate  solution  and  washings  is 
precipitated  with  common  salt,  which  throws  down,  as 
chlorides,  all  the  silver  and  some  of  the  lead  which  may  be 
present  in  the  gold  of  certain  processes.  But  lead  chlo- 
ride is  easily  soluble  in  hot  water  and  is  so  washed  out, 
leaving  pure  silver  chloride.  When  this  is  sufficiently 
washed  it  is  dried  and  melted  with  dry  sodium  carbonate, 
which  decomposes  the  chloride  leaving  pure  metallic  silver 
ready  for  casting  in  bars. 

Another  method  very  commonly  employed  now  is  part- 
ing with  sulphuric  acid,  and  this  method  is  usually  followed 
in  separating  the  silver  and  gold  in  the  cupelled  metal,  de- 
scribed under  silver.  This  silver-gold  alloy  is  boiled  up  in 
iron  kettles  with  strong  sulphuric  acid,  which  dissolves  the 
silver  and  traces  of  other  metals  as  sulphates,  leaving  the 


308  GENERAL  CHEMISTRY. 

gold  as  a  brown,  spongy  mass.  This  is  separated  by 
straining,  washed  thoroughly  with  water,  dried  and  melted 
in  a  crucible  with  sodium  carbonate  and  charcoal,  which 
leaves  it  practically  chemically  pure.  The  solution  of  sil- 
ver sulphate  is  run  into  copper  lined  tanks,  in  which  are 
suspended  sheets  of  copper.  This  metal  displaces  the 
silver,  which  precipitates  as  a  fine  powder,  leaving  copper 
sulphate  as  the  by  product  in  solution.  The  fine  silver  is 
collected  from  the  bottom  of  the  tanks,  thoroughly  washed, 
and  pressed  dry,and  then  fused  with  sodium  carbonate  and  a 
little  charcoal,  yielding  a  product  of  a  high  degree  of  purity. 
The  copper  sulphate  it  recovered  by  crystallization. 

Properties.  Gold  is  the  most  malleable  and  ductile  of 
metals.  It  may  be  beaten  into  foil  so  thin  that  light 
passes  through  it.  Gold  alloys  with  many  heavy  metals 
readily,  as  already  shown.  It  does  not  dissolve  in  any  one 
of  the  common  acids,  but  rather  readily  in  aqua  regia. 

Uses.  The  uses  of  the  metal  are  numerous  and  im- 
portant. It  constitutes  the  standard  coinage  of  most 
nations,  for  this  purpose  being  alloyed  usually  with  copper. 
The  coinage  of  the  United  States  and  France  consists  of 
90  parts  of  gold  to  10  of  copper;  that  of  Great  Britain  is 
a  little  richer  in  gold,  91.666  per  cent.  For  articles  of  jew- 
elry and  ornamentation  the  amount  of  gold  varies  between 
40  and  75  per  cent,  the  fineness  being  usually  expressed  in 
carats.  Pure  gold  is  said  to  be  24  carats  fine;  18  carat 
gold  contains  18  parts  in  24,  or  75  per  cent  of  pure  gold 
and  25  per  cent  of  alloy.  It  is  necessary  to  make  these 
additions  to  gold,  as  alone  it  is  too  soft  for  ordinary  uses. 
Fourteen  carat  gold  preserves  its  bright  color  even  in  the 
worst  laboratory  atmosphere,  and  is  commonly  used  in 
making  rings,  watch  chains  and  cases.  Because  of  its 
great  malleability  gold  is  largely  used  by  dentists  in  the 
filling  of  teeth,  and  is  commonly  employed  in  the  condition 
of  thin  leaf  or  foil.  This  leaf  is  also  used  in  the  lettering 
of  books  and  ornamentation  of  leather  for  many  purposes. 
Gold  is  readily  deposited  as  a  plate  from  a  cyanide  bath 
as  described  under  silver,  and  this  application  is  an  ex- 
tended one. 


GENERAL  CHEMISTRY.  309 

Compounds  Of  Gold,  The  compounds  of  gold  are 
much  less  numerous  than  are  those  of  most  of  the  other 
metals.  Among  the  soluble  salts  the  chloride,  AuCl3,  is 
the  most  important.  It  is  employed  by  photographers  in 
a  toning  solution.  This  chloride  combines  with  the  alkali 
chlorides,  forming  crystalline  double  salts. 

Recognition.  Compounds  of  gold  are  easily  recog- 
nized by  the  dark  brown  precipitate  which  their  solutions 
yield  when  treated  with  a  solution  of  ferrous  sulphate,  also 
by  the  formation  of  a  deep  purple  color  obtained  by 
adding  to  a  gold  solution  a  mixture  of  stannous  and 
stannic  chlorides.  This  color  is  known  as  the  purple  of 
Cassius. 


CHAPTER  XV. 


THE    ALKALNEARTH    GROUP:    BERYLLIUM,    MAG- 
NESIUM, CALCIUM,   STRONTIUM    AND    BA- 
RIUM  THE  SPECTROSCOPE. 

GENERAL  CHARACTERISTICS. 

WE  HAVE  here  a  group  of  five  members  in  which 
certain  marked  resemblances  are  easily  seen.  In 
their  physical  behavior  the  metals  are  much  alike,  with  a 
gradual  change  of  properties,  however,  from  the  lightest, 
beryllium,  to  the  heaviest,  barium.  In  their  compounds 
they  commonly  act  as  bivalent,  forming  oxides,  MO,  and 
hydroxides,  MO2H2.  The  hydroxides  behave  as  strong 
bases.  The  phosphates  and  carbonates  are  insoluble  in 
water,  in  which  respect  they  differ  from  the  alkalies  and 
resemble  the  heavier  metals.  The  sulphates  of  beryllium 
and  magnesium  are  soluble  in  water,  resembling  zinc  sul- 
phate, while  the  sulphates  of  calcium,  strontium  and  barium 
are  characterized  by  their  very  slight  solubility.  The  pro- 
nounced basic  character  of  the  members  of  this  group  in- 
creases with  the  atomic  weight.  Beryllium  does  not  de- 
compose water  at  any  temperature, magnesium  acts  slowly 
on  hot  water,  while  barium  behaves  as  energetically  as 
sodium  and  potassium.  Beryllium  and  magnesium  hydrox- 
ides are  practically  insoluble  in  water  and  both  decompose 
at  comparatively  low  temperatures,  leaving  oxides;  the 
other  hydroxides  are  more  soluble  and  more  stable;  in  fact, 
barium  hydroxide  can  be  fused  like  KOH  or  NaOH, 
and  it  is  soluble  enough  to  form  a  strongly  alkaline  solu- 
tion. The  following  table  shows  the  range  of  atomic  vol- 
umes and  other  properties : 


GENERAL  CHEMISTRY. 


311 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

BOILING 
POINT. 

Beryllium  .... 
Magnesium  ..  . 
Calcium. 

9.08 
24.28 
40  07 

1.64 

1.75 
1  58 

55 
13.9 
25  4 

940° 

750° 

Very  high 
1100° 

Strontium 

87.61 

2  54 

34  5 

Red  heat 

Barium  .... 

187.43 

3.75 

36.6 

475° 

The  melting  and  boiling  points  in  this  group  have  not 
been  very  accurately  determined.  The  values  given  are 
probably  close  approximations,  but  here,  as  in  the  alkali 
group,  we  notice  an  increase  in  atomic  volume  and  decrease 
in  melting  point  with  an  increase  in  the  atomic  weight; 
conditions  which  must  have  some  rntimate  connection 
with  the  nature  of  the  atoms  themselves. 


BERYLLIUM. 

Occurrence.  This  metal  is  found  in  a  number  of  com- 
paratively rare  minerals,  of  which  the  most  important  are 
beryl  and  emerald,  which  are  essentially  Be3Al2(SiO3)6, 
and  chrysoberyl,  BeAl2O4.  The  term  emerald  is  applied  to 
transparent  beryls  with  a  green  tinge,  while  if  they  have  a 
blue  color  they  are  called  aquamarins. 

History.  The  oxide  was  recognized  in  1797  and  the 
metal  was  isolated  in  1828  by  fusion  of  the  chloride  with 
metallic  potassium.  It  has  been  isolated  by  several  other 
processes  since,  and  has  been  found  to  be  somewhat  mal- 
leable and  capable  of  taking  a  polish. 

Compounds  of  Beryllium.  These  are  of  little  impor- 
tance. The  salts  are  mostly  soluble  in  water  and  possess  a 
bitter  sweetish  taste,  by  which  they  may  be  recognized. 
The  chloride  and  sulphate  are  easily  obtained  in  crystal- 
line condition.  On  account  of  the  sweetish  taste  the  metal 
is  sometimes  called  glucinum. 


312  GENERAL  CHEMISTRY. 

MAGNESIUM. 

Occurrence.  It  is  found  in  a  large  number  of  miner- 
als and  very  widely  distributed.  The  best  known  are  the 
carbonate  or  magnesite,  MgCO3,  dolomite,  MgCO3.CaCO3, 
kieserite,  MgSO4.H2O,  carnallite,  MgCl2.KC1.6H2O,  and 
several  silicates.  The  sulphate  and  chloride  are  found  in 
many  mineral  springs. 

History.  The  sulphate  was  produced  from  the  water 
of  Epsom  springs  in  the  seventeenth  century,  and  the  oxide 
or  magnesia  alba  obtained  from  it  in  1707.  As  this  sub- 
stance was  first  made  it  was  very  impure  and  was  for  a  long 
time  confounded  with  lime,  but  Black  showed  the  difference 
between  them  in  1755.  Davy  attempted  to  isolate  the 
metal  which  he  recognized  the  magnesia  alba  must  con- 
tain, but  was  not  able  to  secure  it  in  pure  form.  This  was 
accomplished  by  Bussy  in  1830,  who  fused  the  dry  chlo- 
ride with  potassium: 

MgCl8  +  K8  =  Mg+2KCl. 

Preparation.  The  method  just  mentioned  served  for 
the  preparation  of  the  metal  for  years.  Bunsen  showed  in 
1852  that  it  could  be  isolated  by  electrolysis  of  the  fused 
chloride,  and  methods  based  on  this  behavior  are  in  use 
on  the  large  scale.  A  great  deal  of  magnesium  is  made  by 
fusing  a  mixture  of  the  dry  chloride  with  fluorspar  and 
metallic  sodium,  the  latter  being  cheaper  than  potassium. 
The  reduced  metal  is  now  purified  by  distillation. 

Properties.  Magnesium  is  a  white,  malleable  and 
ductile  metal  which  may  be  drawn  into  wire  when  warm. 
It  breaks  with  a  crystalline  fracture.  When  heated  in  the 
air  to  a  high  temperature  it  burns  with  an  intense  white 
light. 

Ex.  162.  Heat  a  short  piece  of  magnesium  wire  in  the  flame  of  a 
Bunsen  burner.  When  combustion  begins  remove  it  from  the  lamp 
and  allow  it  to  burn  freely  in  the  air.  The  light  is  often  employed  in 
photography. 

Uses.     As  just  suggested,  in  photography,  as  the  light 


GENERAL  CHEMISTRY.  313 

is  strongly  actinic.  The  powder,  pure  or  mixed  with  potas- 
sium chlorate,  is  generally  used  instead  of  the  wire,  because 
the  combustion  is  more  intense. 

Magnesium  Oxide.  This  substance,  known  also  as 
magnesia,  is  the  white  powder  which  is  formed  when  the 
metal  is  burned  in  air  or  in  oxygen.  Practically  it  is  made 
by  strongly  heating  the  carbonate,  which  decomposes  as 
follows: 


The  oxide  is  practically  insoluble  in  water.  Magnesium 
hydroxide,  MgO2H2,  is  formed  by  precipitating  magnesium 
sulphate  or  chloride  by  sodium  hydroxide,  or  by  long  con- 
tact of  the  oxide  with  water. 

Magnesium  Carbonate  is  found  in  nature  in  nearly 
pure  condition,  and  mixed  with  calcium  carbonate,  as 
dolomite.  An  artificial  product  is  easily  made  by  precipi- 
tating the  sulphate  by  sodium  carbonate,  but  contains 
always  some  magnesium  hydroxide.  In  presence  of  ammo- 
nium salts  precipitation  is  very  imperfect.  Ammonium 
carbonate  only  partially  precipitates  magnesium  salts.  If 
the  solution  contains  a  sufficient  amount  of  ammonium 
chloride  no  precipitate  whatever  forms. 

Ex.  163.  Let  the  student  prepare  a  solution  of  magnesium  sul- 
phate and  divide  it  into  several  portions.  To  one  add  sodium  carbonate, 
to  another  ammonium  carbonate  and  to  a  third  and  fourth  portions 
ammonium  chloride  solution  and  then  these  carbonates.  Observe  the 
different  behaviors. 

The  following  equation  represents  the  reaction  between 
magnesium  sulphate  and  sodium  carbonate: 

5MgSO4-f  5Na2CO3+6H2O  = 

5Na2SO4+4MgCO3.MgO2H2.5H2O-f-CO2. 

The  composition  of  the  precipitate  varies  with  the  con- 
centration and  temperature  of  the  solutions. 

Magnesium  Chloride  is  a  soluble  crystalline  salt 
obtained  by  dissolving  the  carbonate  in  hydrochloric  acid. 
It  is  used  in  analytical  chemistry. 


314  GENERAL  CHEMISTRY. 

Magnesium  Sulphate.  This  salt  has  been  already 
referred  to.  It  is  made  by  purifying  the  native  mineral  or 
by  the  action  of  sulphuric  acid  on  dolomite.  The  sulphates 
of  magnesium  and  calcium  are  formed,  but  the  latter  is  so 
slightly  soluble  in  water  that  a  separation  is  easily  made, 
permitting  the  magnesium  compound  to  be  obtained  in 
practically  pure  condition.  It  has  many  uses,  and  is  the 
starting  point  in  the  manufacture  of  other  compounds  of 
the  metal.  As  commonly  met  with  in  commerce  it  is  in  the 
crystalline  condition  known  as  Epsom  salt,  MgSO4.7H2O. 
It  will  be  seen  later  that  we  have  here  the  first  one  of  an 
important  group  of  sulphates  which  crystallize  with  7 
molecules  of  water  and  which  are  isomorphous  and  similar 
in  many  respects. 

Magnesium  Phosphates.  Several  are  known,  distin- 
guished by  their  insolubility.  Ammonium  magnesium 
phosphate,  or  triple  phosphate,  NH4MgPO4.6H2O,  is  a 
characteristic  crystalline  precipitate. 

Recognition.  Magnesium  compounds  may  be  recog- 
nized by  the  formation  of  the  precipitate  just  mentioned, 
obtained  by  adding  sodium  phosphate  to  a  magnesium  so- 
lution containing  ammonium  chloride  and  hydroxide. 


CALCIUM. 

Occurrence.  The  metal  is  found  as  carbonate  in  lime- 
stone, marble,  chalk  and  coral ;  as  a  sulphate  in  gypsum 
and  anhydrite  ;  as  phosphate  in  bone  ash  and  in  apatite  ; 
as  silicate  in  many  minerals  and  less  abundantly  in  other 
combinations. 

History  and  Preparation.  Lime  was  known  to  the 
ancients,  and  until  the  beginning  of  this  century  was  sup- 
posed to  be  a  simple  substance.  Davy  showed  that  this  is 
the  oxide  of  the  metal  which  he  succeeded  in  isolating  by 
electrolysis  of  the  chloride.  It  has  never  been  separated 
in  large  quantity  and  the  small  amounts  obtained  by  chem- 
ists seldom  weigh  more  than  a  few  grams. 


GENERAL  CHEMISTRY.  315 

Properties  and  Uses.  Calcium  is  a  yellow  metal 
much  like  brass  in  appearance,  which  oxidizes  readily  in 
moist  air  and  decomposes  water  quickly.  It  may  be  pre- 
served under  benzine.  When  ignited  it  burns  with  a  bright 
yellow  light.  Because  of  the  expense  of  isolating  it,  it  has 
no  technical  applications. 

Calcium  Oxide.  This  substance  is  known  commonly 
as  limet  having  the  formula  CaO.  It  is  made  by  strongly 
heating  limestone  or  marble  : 

CaCO3  =  CaO+CO2. 

This  decomposition  is  usually  carried  out  in  lime  kilns 
with  the  loss  of  the  CO2. 

Calcium  Hydroxide.  This  substance  is  known  as 
slaked  lime,  and  has  the  composition  represented  by  the 
symbols  CaO2H2  or  Ca(OH)2. 

Ex.  164.  Put  a  piece  of  soft,  well  "burned  lime,  as  large  as  a  wal- 
nut, in  a  porcelain  dish,  and  pour  over  it  some  lukewarm  water,  a  few 
drops  at  a  time.  The  lime  absorbs  the  water  and  soon  begins  to  swell 
and  then  crumble.  After  the  piece  has  become  thoroughly  disintegrated 
add  enough  water  to  make  a  thick  liquid  and  stir  well  with  a  glass  rod. 
The  product  is  known  as  milk  of  lime.  Pour  a  few  drops  into  a  test-tube 
and  add  some  more  water,  and  then  phenol-phthalein.  A  red  color  is 
produced,  showing  that  the  substance  is  an  alkali.  Into  another  test- 
tube  pour  more  of  the  milky  liquid,  add  an  equal  volume  of  water  and 
then  filter  through  a  small  filter  into  a  clean  flask  or  beaker.  Add  phe- 
nol-phthalein to  the  filtrate.  It  is  found  to  be  alkaline.  Then  make 
some  very  dilute  hydrochloric  acid,  by  diluting  the  weakest  labora- 
tory acid  with  ten  times  the  volume  of  water  (1  Cc.  with  10  of  water  will 
be  enough),  and  pour  a  little  of  this  into  the  colored  filtrate.  Notice 
that  the  color  is  soon  discharged,  showing  that  the  amount  of  real  alkali 
present  must  be  very  small.  The  slaked  lime  is  only  slightly  soluble  in 
water.  Now  divide  what  remains  of  the  milk  of  lime  into  two  poitions; 
to  one  add  strong  hydrochloric  acid  and  to  the  other  add  strong  nitric 
acid.  Clear,  or  nearly  clear  solutions  are  obtained,  showing  that  the 
slaked  lime,  while  but  slightly  soluble  in  water,  is  readily  soluble  in  the 
two  acids,  but  with  formation  of  new  substances. 

The  clear  solution  of  the  slaked  lime  in  water  is  known 
as  lime-water.  About  700  parts  of  cold  water  are  required 
to  dissolve  one  part  of  the  substance.  It  is  much  less  sol- 


316  GENERAL  CHEMISTRY. 

uble   in    hot    water.     The  reaction  which   takes  place  in 
slaking  the  lime  is  represented  by  the  equation  : 


With  a  moderate  amount  of  water  this  slaked  lime 
yields  milk  of  time,  while  in  a  great  excess  it  is  soluble, 
yielding  lime-water.  The  behavior  of  the  clear  solution 
with  carbonic  acid  gas  from  the  air  or  from  the  lungs  has 
been  shown  already.  This  lime-water  is  therefore  a  very 
useful  test  for  CO2,  and  is  commonly  used  for  the  pur- 
pose. 

Slaked  lime  is  used  mainly  in  making  mortar,  but  also 
in  great  quantities  in  chemical  industries,  as  illustrated  in 
the  liberation  of  ammonia. 

Common  mortar  is  a  mixture  of  slaked  lime  and  sand 
and  hardens  mainly  through  the  absorption  of  carbon  di- 
oxide by  the  lime,  thus  forming  carbonate. 

Calcium  Carbonate.  As  stated  above,  this  occurs  in 
nature  in  various  forms,  and  may  be  produced  by  precipi- 
tating a  solution  of  calcium  chloride  with  one  of  ammonium 
or  sodium  carbonate.  In  this  form  it  is  a  fine  white 
powder  and  has  several  uses.  Calcium  carbonate  is  very 
nearly  insoluble  in  pure  water,  but  dissolves  appreciably 
in  water  containing  carbonic  acid,  with  formation  of  bicar- 
bonate : 

CaCO3+CO2+H2O  =  H2Ca(CO3)2. 

The  temporary  hardness  of  most  natural  water  is  due 
to  the  presence  of  this  and  the  analogous  magnesium 
bicarbonate.  Both  compounds  are  destroyed  by  boiling 
the  water. 

Calcium  Chloride.  This  salt  is  usually  made  by  dis- 
solving marble  in  hydrochloric  acid.  It  comes  into  com- 
merce as  the  anhydrous  salt,  CaCl2,  or  a  crystallized 
product,  CaCl2.6H2O.  A  partially  dried  product  with 
2H2O  is  obtained  from  the  latter  by  heating  to  200°,  and  is 
used  often  instead  of  the  anhydrous  salt  in  drying  gases  or 
organic  liquids.  Calcium  chloride  is  a  by-product  in 
several  chemical  industries. 


GENERAL  CHEMISTRY,  317 

Calcium  Sulphate.  This  important  substance  is 
represented  by  the  formula  CaSO4.  In  nature  it  is  found 
in  large  quantities  as  gypsum,  CaSO4.2H2O.  When  a 
gentle  heat  is  applied  to  the  gypsum  it  loses  its  water  of 
crystallization  and  forms  the  dry  substance,  or  plaster  of 
Paris.  This  plaster  of  Paris  is  distinguished  by  its  prop- 
erty of  hardening  when  mixed  with  a  little  water,  and  is 
therefore  used  as  a  cement,  and  as  a  substance  for  taking 
casts  and  making  surgical  dressings.  The  use  in  making 
casts  or  copies  of  coins  and  medals  can  be  shown  by  the 
following  experiment: 

Ex.  165.  Rub  a  little  oil  over  a  large  coin  or  medal  and  then 
remove  most  of  it  with  the  finger,  leaving  enough,  however,  to  form  a 
very  thin  uniform  layer  over  the  metal  surface.  Then  mix  15  to  20  Gm. 
of  plaster  of  Paris  with  enough  water  to  form  a  thick  cream,  and  without 
delay,  pour  this  over  the  coin  so  as  to  form  a  layer  half  a  centimeter  or 
more  in  thickness.  A  strip  of  paper  should  be  wrapped  around  the 
coin  so  as  to  form  a  cell  into  which  the  creamy  mass  is  poured.  In  a 
few  minutes  the  plaster  "sets,"  but  it  should  be  allowed  to  remain  long 
enough  to  become  quite  hard.  The  solid  plaster  c&st  can  then  be  re- 
moved very  easily  and  shows  every  feature  of  the  coin  in  reversed  po- 
sition. 

When  gypsum  is  heated  to  a  temperature  of  200°  it 
becomes  overburnt  and  will  no  longer  take  up  water  readily. 

Other  Compounds.  Several  of  these  must  be  men- 
tioned briefly.  Calcium  nitrate  is  found  as  a  deposit  in 
many  caves  and  is  called  "cave  niter."  Calcium  hypo- 
chlorite  is  the  active  constituentof  bleachingpowderand  has 
the  formula  CaO3Cl2.  It  is  made  by  passing  chlorine  gas 
over  slaked  lime.  The  bleaching  powder  of  commerce  is 
a  mixture  of  the  hypochlorite,  with  calcium  chloride  result- 
ing in  the  reaction  and  some  unchanged  lime.  Calcium 
phosphate  in  crude  form  is  found  in  bone  ash  from  the 
burning  of  bones,  also  in  large  deposits  in  the  earth,  some 
of  which  have  had  their  origin  in  the  bones  of  animals. 
Calcium  metaborate  is  an  important  mineral,  and  the 
source  of  much  of  the  boric  acid  and  borax  of  commerce. 
The  carbide  of  calcium,  CaC2,  is  now  an  important  sub- 
stance produced  by  subjecting  a  mixture  of  coke  and  lime 
to  the  heat  of  the  electric  arc.  With  water  it  yields  acetylene, 
already  described. 


318  GENERAL  CHEMISTRY. 

Recognition.  Calcium  compounds  are  recognized  by 
the  reddish  yellow  color  they  impart  to  the  flame  of  the 
Bunsen  burner,  by  the  bright  red  and  bright  green  line  they 
exhibit  in  the  spectroscope,  and  by  the  white  precipitates 
produced  in  their  solutions  when  made  alkaline  and  mixed 
with  solutions  of  ammonium  carbonate  oroxalate.  Precip- 
itated calcium  oxalate,  CaC2O4.2H2O,  is  a  very  insoluble 
substance. 

STRONTIUM. 

Occurrence.  The  metal  is  comparatively  rare,  being 
found  in  a  few  minerals  only  and  these  not  abundant.  The 
sulphate,  celestine,  SrSO4,  and  the  carbonate,  strontianite, 
SrCO3,  are  the  most  important. 

History.  The  carbonate  was  first  mistaken  for  barium 
carbonate,  but  in  1792  the  distinction  was  shown.  In  1808 
Davy  isolated  the  metal  in  impure  condition  by  electrol- 
ysis of  the  chloride.  It  is  still  prepared  in  that  way,  but 
not  easily,  and  is  therefore  rare  as  a  metal. 

Properties.  It  is  a  yellow  metal  which  oxidizes  very 
readily  and  decomposes  water  with  great  energy.  It  has 
no  scientific  or  technical  uses. 

Compounds  of  Strontium.  Of  these  the  most  impor- 
tant are  the  oxide,  the  chloride,  the  nitrate  and  the  chlo- 
rate. The  oxide  yields  the  hydroxide,  SrO2H2,  which 
combines  readily  with  sugar,  forming  a  crystalline  saccha- 
rate  of  some  importance  in  sugar  refining.  The  nitrate  and 
chlorate  are  used  mainly  in  fireworks  mixtures,  to  which 
they  impart  a  bright  red  color. 

Recognition.  Strontium  compounds  color  the  flame 
of  the  Bunsen  burner  bright  red,  they  also  show  a  number 
of  characteristic  red  lines  in  the  spectroscope.  Solutions 
of  strontium  compounds  yield  a  white  precipitate  of  the 
carbonate  when  treated  with  ammonium  carbonate. 

BARIUM. 
Occurrence.     This  element  is  found  mainly  in  heavy 


GENERAL  CHEMISTRY,  319 

spar,  BaSO4,  and  in  the  carbonate  or  witherite,  BaCO3,the 
first  of  which  is  a  comparatively  common  mineral. 

History  and  Preparation.  Some  of  the  properties  of 
heavy  spar  have  been  known  since  early  in  the  seventeenth 
century.  Scheele,  in  1774,  showed  that  it  contained  a 
new  earth  and  Davy  showed  that  this  earth  is  the  oxide  of 
the  metal  now  called  barium.  His  attempts  at  isolating 
the  metal  in  pure  condition  were  not  successful;  Bunsen 
and  others  later  succeeded  in  obtaining  it  by  electrolysis, 
but  the  process  is  a  difficult  one. 

Properties.  The  metal  appears  to  be  lighter  colored 
than  calcium  and  strontium;  but  it  oxidizes  with  extreme 
readiness,  so  that  it  is  probable  that  a  pure  substance  has 
never  been  secured.  Authorities  differ  as  to  the  melting 
point  of  the  metal.  By  some  it  is  placed  higher  than  that 
of  cast  iron,  while  by  others  it  is  given  as  lower  than  that 
of  strontium.  The  metal  has  no  uses  in  the  arts. 

Barium  Oxide,  BaO,  is  obtained  by  ignition  of  the 
nitrate  or  carbonate  at  a  high  temperature.  To  reduce 
the  carbonate  in  quantity  it  is  best  to  add  carbon  to  it. 
The  pure  oxide  is  a  heavy  white  powder,  which  when 
heated  in  the  air  or  oxygen  to  redness  absorbs  a  second 
atom  of  oxygen  with  formation  of  the  dioxide,  BaO2.  This 
second  atom  is  given  off  at  a  bright  red  heat  and  on  this 
behavior  is  founded  a  method  of  preparing  oxygen  from 
the  air  in  quantity.  The  dioxide  is  used,  as  already  illus- 
trated, in  the  production  of  hydrogen  dioxide. 

Barium  Hydroxide.  The  oxide,  BaO,  combines  very 
readily  with  water  forming  BaO2H2,  which  dissolves  in 
considerable  quantity  and  forms  a  crystalline  product, 
BaO2H2.8H2O.  This  is  a  valuable  reagent  because  of  its 
strong  alkaline  reaction  and  power  of  absorbing  CO2.  It 
is  also  used  in  the  separation  of  sugar  from  solutions. 

Barium  Chloride.  This  salt  is  prepared  by  dissolving 
barium  carbonate  in  hydrochloric  acid.  It  is  readily  solu- 
ble in  water,  from  which  it  crystallizes  as  BaCl2.2H8O.  A 


320  GENERAL  CHEMISTRY. 

solution  of  this  salt  is  a  valuable  reagent  for  the  precipita- 
tion of  sulphates. 

Other  Compounds.  The  nitrate,  Ba(NO3)2,  and  the 
chlorate,  Ba(ClO3)2,  are  used  in  fireworks  to  impart  a 
greenish  yellow  color.  The  sulphate,  in  the  form  of  ground 
heavy  spar,  is  very  commonly  used  as  an  adulterant  in  white 
paints.  It  has  no  value  as  a  pigment  itself. 

Recognition.  Barium  compounds  color  the  Bunsen 
flame  greenish  yellow,  and  in  the  spectroscope  show  a 
number  of  yellow  and  green  lines.  Sulphates  produce  in 
barium  solutions  a  heavy  white  precipitate  of  barium  sul- 
phate, which  is  extremely  insoluble  in  water  and  acids. 


THE  USE  OF  THE  SPECTROSCOPE. 

It  has  been  shown  already  that  several  metals  may 
be  detected  by  the  colors  which  their  salts  impart  to 
the  flame  of  the  Bunsen  burner,  in  which  they  are  heated 
on  a  platinum  wire.  When  a  clean  platinum  wire  is  dipped 
in  a  strong  solution  of  the  chloride  or  nitrate  of  calcium, 
barium,  strontium,  sodium,  potassium  or  lithium  and  then 
held  in  a  Bunsen  flame  a  characteristic  color  is  produced 
which  is  sufficient  for  the  immediate  identification  of  the 
metal,  as  long  as  one  only  is  present.  Suppose,  however, 
that  we  have  a  potassium  salt  mixed  with  an  equal  or 
greater  quantity  of  a  salt  of  lithium,  sodium,  calcium  or 
strontium.  It  will  be  found  now  on  making  the  test  that 
the  potassium  color  is  wholly  obscured  or  uncertain.  In 
the  same  manner  it  may  be  shown  that  a  small  amount  of 
a  calcium  salt  may  be  hidden  by  strontium,  sodium,  or 
barium  even.  The  application  of  these  flame  tests  to  mix- 
tures may  lead,  therefore,  to  quite  imperfect  conclusions, 
because  the  eye  is  not  able  to  recognize  certain  colors  in 
presence  of  others. 

Now,  the  flame  appears  colored  by  the  volatilization  of 
these  salts,  because  at  a  high  temperature  they  are  decom- 
posed and  the  products  of  decomposition  thrown  into  very 
rapid  rates  of  vibration.  These  different  rates  of  vibration 


GENERAL  CHEMISTRY. 


321 


are  in  turn  communicated  to  the  surrounding  ether,  and 
this  is  the  medium  which  brings  the  impression  to  the  eye. 
What  is  called  light  is,  outside  the  body,  a  rate  of  motion. 
Different  rates  of  ether  motion  conveyed  through  the  trans- 
parent media  of  the  eye  to  the  retina  produce  there  differ- 
ent impressions,  which,  communicated  by  the  optic  nerve 
to  the  brain,  give  to  the  individual  the  sensations  of  differ- 
ent kinds  of  light.  These  sensations  are  commonly  result- 
ant effects  produced  by  the  blending  of  different  simul- 
taneous disturbances  of  the  same  part  of  the  retina. 

Suppose,  however,  that  the  light  from  the  colored  flame 


FIG.    27, 

can  reach  the  eye  only  after  passing  through  a  colorless 
transparent  prism,  of  glass  preferably.  Two  new  phenom- 
ena are  immediately  noticed.  The  light  is  deviated  or 
bent  from  its  direct  course;  that  is,  it  has  suffered  refrac- 
tion, and  it  has  also  been  dispersed  ox  broken  up  into  lights 
of  various  shades.  This  is  due  to  the  fact  that  the  differ- 
ent ethereal  disturbances  are  conveyed  through  the  sub- 
stance of  the  glass  with  different  degrees  of  facility.  The 
light  motions  in  some  cases  are  more  retarded  than  in 
others,  and  hence  the  separation  of  the  component  parts  of 
what  is  termed  the  beam  of  light  after  it  leaves  the  prism. 
The  eye  brought  opposite  different  points  of  the  prism,  in 
the  same  plane  perpendicular  to  the  refractive  surfaces  of 
the  prism,  receives  now  separate  and  distinct  sensations 


322 


GENERAL  CHEMIS  TK  J \ 


as  it  would  from   several  colored  lights  burning   near  to- 
gether, one  blue,  one  red,  one  green,  and  so  on. 

The  images  reaching  the  eye  are  rendered  sharper  and 
more  clearly  distinguishable  if  a  very  narrow  vertical  slit 
is  placed  in  front  of  the  prism,  and  if  between  the  latter 
and  the  eye  we  have  a  double  convex  lens  or  small  tele- 
scope to  take  up  the  rays  and  throw  them  into  the  eye  in 
nearly  parallel  direction.  Such  an  arrangement  would 


FIG.  28. 

constitute,  in  fact,  a  simple  spectroscope,  and  is  shown  dia- 
grammatically  in  Fig.  27. 

The  actual  arrangement  of  a  spectroscope  is  shown  in 
Fig.  28. 

The  optical  parts  of  the  instrument  are  supported  on 
the  base,  F.  The  prism,  P,  has  a  refractive  angle  of  60°, 
usually.  B  is  a  tube  furnished  with  lenses  and  constitutes 
an  observing  telescope.  C  is  the  collimator  tube  which 
receives  and  renders  parallel  the  rays  of  light  coming  to  the 
prism  from  the  flame  or  other  object  under  examination. 
It  has  a  double  convex  lens  in  the  end  near  the  prism,  and 


GENERAL  CHEMISTRY.  323 

at  the  other  end  is  the  slit,  which  can  be  adjusted  so  as  to 
be  exactly  in  the  focus  of  the  lens.  By  this  arrangement 
the  light  rays  thrown  on  the  prism  are  made  nearly  parallel. 
The  tube  A  carries  at  its  outer  end  a  fine  scale  photo- 
graphed on  glass.  This  can  be  illuminated  by  a  lamp  and 
its  image  so  thrown  on  the  surface  of  P  as  to  be  reflected 
into  the  telescope,  B.  The  lines  on  this  scale  serve  to 
designate,  arbitrarily,  the  position  of  the  colored  bands  or 
lines  of  the  spectrum  produced  and  seen  at  the  same  time. 

In  Fig.  29  we  have  an  illustration  of  the  slit  through 
which  the  light  enters  the  instrument. 

This  slit  must  always  be  very  narrow,  but  its  width  may 
be  varied  by  the  screw  shown  at  the  right.  The  lower  half 
of  the  slit  is  covered  with  a  small  prism.  Light  enters  the 
upper  part  of  the  slit  from  directly  in  front,  while  byreflec- 


FIG.  89. 

tion  from  one  of  the  surfaces  of  the  small  prism  it  may  be 
made  to  enter  the  lower  part  of  the  slit  from  the  side.  By 
this  arrangement  it  is  possible  to  compare  two  different 
lights  and  obtain  two  different  spectra,  one  above  the 
other.  In  the  illustration  the  lamp,  D,  is  in  front  of  the 
slit,  while  E  is  at  one  side  and  furnishes  light  to  the 
prism,  P,  by  reflection  from  the  small  prism  on  the  slit 
front. 

The  working  of  the  apparatus  as  a  whole  may  now  be 
illustrated.  If  we  place  in  front  of  the  open  slit  an  ordi- 
nary illuminating  gas  burner,  or  if  we  throw  sunlight 
directly  into  it,  the  eye  at  the  end  of  B  receives  the  ordi- 
nary solar  or  continuous  spectrum.  If,  however,  a  Bunsen 
burner  be  placed  in  front,  and  in  this  a  little  sodium  chlo- 
ride be  volatilized,  the  eye  perceives  now  a  single  bright 


324 


GENERAL  CHEMISTRY. 


yellow  band,  which  is  merely  the  image  of  the  slit,  and 
varies   in  width  with  the  width  of  the  latter.     If  calcium 


Red. 


Violet. 


D  ,         I  Or-  I  Yel-  I       c 
Red'        |ange.|low.  |       Green' 


Bine. 


Violet. 


FIG.  30. 


Showing  the  number  and  positions  of  the  important  lines  in  the 
spectra  of  some  of  the  metals. 

chloride  be  volatilized  in  theflame  several  lines  will  be  seen, 
two  of    which,  a  bright  green  and  a  bright    red,  are  promi- 


GENERAL  CHEMISTRY.  325 

nent.  A  salt  of  strontium  yields  a  number  of  bright  lines, 
mostly  red.  Further  examination  will  show  that  the 
results,  as  far  as  the  positions  of  the  lines  are  concerned, 
are  the  same  whatever  salts  of  the  metals  we  use  and 
whatever  the  temperature,  within  rather  wide  limits.  Our 
ordinary  spectrum  analysis  is  based  on  these  facts. 

By  use  of  the  illuminated  scale  we  may  note  the  rela- 
tive positions  of  all  these  lines,  and  make  a  chart  of  them 
such  as  is  shown  in  the  above  figure.  The  color  of  the  lines 
is  shown  by  the  colors  of  the  solar  spectrum  designated  in 
the  upper  portion  of  the  figure.  The  positions  of  the  lines 
vary  on  the  scale  with  the  nature  of  the  prism.  Hence,  a 
chart  made  in  this  manner  can  show  the  results  obtained  in 
a  particular  instrument  only. 

Observations  are  best  made  in  a  dark  room,  the  light 
entering  the  apparatus  from  the  flame  under  examination 
only.  But  for  the  ordinary  purposes  very  good  results  can 
be  secured  by  mounting  the  spectroscope  in  a  dark  corner 
of  a  laboratory,  protected  from  acid  and  hydrogen  sulphide 
fumes.  The  prism  should  be  covered  with  a  dark  paper 
cap,  with  openings  opposite  the  tubes  only. 

While  all  salts  of  a  metal  yield  at  a  temperature  suffi- 
ciently elevated  the  same  bands,  for  practical  purposes  it  is 
best  to  work  with  those  easily  volatilized,  as  the  chlorides 
or  nitrates.  The  intensity  of  the  spectrum  is  usually 
increased  by  moistening  the  substance  under  examination 
with  hydrochloric  acid.  Some  of  the  semi-solid  mixture  or 
strong  solution  is  taken  up  on  the  loop  of  a  platinum  wire 
and  held  in  the  flame,  as  shown  in  the  illustration. 

With  a  little  practice  the  student  can  make  himself 
thoroughly  familiar  with  the  number  and  position  of  the 
lines  given  by  the  different  metals.  With  these  in  mind 
he  is  able  to  control,  to  some  extent,  the  results  of  his 
analyses  by  precipitation  and  to  detect  the  presence  of 
several  elements  briefly  discussed  in  the  preceding  pages. 
Lithium,  for  example,  is  a  constituent  of  many  natural 
waters  and  it  can  be  most  readily  found  by  concentrating 
the  water  to  a  small  volume  and  applying  the  test  to  the 
residue.  For  small  traces  of  lithium,  with  large  quantities 
of  other  salts,  the  latter  must  be  removed  by  precipitation 


326  GENERAL  CHEMISTRY. 

before  making  the  actual  test,  as  will  be  suggested  below. 
By  the  aid  of  the  spectroscope  several  metals  have  been 
discovered.  Of  these  may  be  mentioned  caesium  and 
rubidium,  by  Bunsen  and  Kirchhoff;  thallium,  by  Crookes; 
and  gallium,  by  Lecoq  de  Boisbaudran.  As  an  illustration 
of  the  detection  of  lithium,  let  the  student  evaporate  sev- 
eral hundred  cubic  centimeters  of  artesian  water  to  dry- 
ness  in  a  porcelain  dish.  To  the  evaporating  solution  add 
a  feiu  drops  of  solution  of  pure  potassium  carbonate.  Heat 
the  residue  strongly,  allow  it  to  cool  and  then  add  a  few 
Cc.  of  distilled  water,  boil  and  filter.  Wash  the  residue 
with  1  Cc.  of  boiling  water  and  add  this  filtrate  to  the  other. 
By  this  treatment  much  of  the  calcium,  barium,  stron- 
tium, magnesium,  iron  and  other  metals  possibly  present,  is 


FIG.  31. 

separated.  Next  evaporate  the  filtrate  nearly  to  dryness, 
after  adding  enough  hydrochloric  acid  to  change  the  reac- 
tion. Mix  the  residue  with  strong  alcohol,  using  a  few  Cc. 
only.  Stir  and  filter.  Lithium  chloride,  being  soluble  in 
alcohol,  is  separated  in  this  manner  from  the  bulk  of  the 
other  alkali  chlorides.  Evaporate  the  alcohol  filtrate  and 
test  the  residue  with  the  spectroscope.  A  single  bright 
red  line  is  characteristic  of  lithium,  but  in  this  and  nearly 
all  other  tests  the  yellow  line  of  sodium  will  be  seen 
because  of  the  practical  difficulty  of  separating  traces  of 
this  metal  from  substances  under  examination,  and  because, 
further,  of  the  very  great  delicacy  of  the  sodium  reaction,  a 
minute  trace  being  sufficient  to  show  a  sharp  line. 

Absorption  Spectra. 

To  illustrate  another  very   important  branch  of  spec- 
trum analysis  let  the  student  prepare  very  dilute  solutions 


GENERAL  CHEMISTRY.  327 

of  potassium  dichromate,  potassium  permanganate,  ani- 
lin  red  and  other  substances  which  yield  colored  solu- 
tions. Pour  these  in  test-tubes,  which  support  between  the 
slit  of  the  spectroscope  and  a  luminous  gas  or  oil  lamp 
flame.  It  will  be  observed  that  a  part  only  of  the  continu- 
ous spectrum  is  now  visible,  some  of  the  colors  being  ab- 
sorbed by  the  solutions  in  the  tubes.  The  position  and 
number  of  these  zones  or  bands  of  absorption  are  charac- 
teristic for  solutions  of  many  colored  substances.  Appli- 
cations of  these  principles  are  found  mainly  in  the  exam- 
ination of  organic  coloring  matters. 

Direct  Vision  Spectroscopes. 

For  many  purposes  much  simpler  apparatus  than  de- 
scribed may  be  employed  in  the  examination  of  metallic 
spectra.  The  direct  vision  spectroscope  is  an  instrument 
illustrated  in  the  last  figure  in  which  the  flame,  slit,  prisms 
and  eye  are  in  one  straight  line.  The  dimensions  of  the 
apparatus  may  be  thus  much  reduced.  By  a  combination 
of  small  prisms,  mounted  in  a  brass  tube,  it  is  possible  to 
correct  the  refraction,  without  eliminating  their  dispersion 
effect.  Direct  vision  spectroscopes  are  usually  made 
small  and  are  employed  for  special  purposes  only.  But 
occasionally  they  are  made  with  large  prisms  for  more 
elaborate  investigations.  In  Fig.  31  the  light  enters 
through  A  and  emerges  at  C. 


CHAPTER  XVI. 


ZINC,  CADMIUM  AND  MERCURY. 


GENERAL  CHARACTERISTICS. 

THESE  elements  bear  the  same  relation  to  the  alkali- 
earth  metals  that  copper,  silver  and  gold  bear  to  the 
true  alkali  metals.  In  all  of  their  important  compounds 
zinc  and  cadmium  are  bivalent,  while  mercury  forms  two 
series  of  compounds,  the  mercurous  and  mercuric.  The 
latter  are  the  most  common  and  stable.  The  table  below 
exhibits  the  important  physical  characteristics  of  the  three 
metals: 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

BOILING 
POINT. 

VAPOR 
DENSITY. 

H  =  l 

Zinc  

65.41 

7.15 

9.15 

420° 

930° 

34.3 

Cadmium  

111.95 

8.05 

12.94 

320 

770 

50,8 

Mercury  

200.00 

13.59 

14.71 

—39 

357 

98.2 

It  will  be  observed  that  here,  as  before,  we  have  a 
decrease  in  the  melting  and  boiling  points  corresponding  to 
an  increase  in  the  atomic  weight.  The  vapor  densities  of 
these  metals  have  been  determined  with  considerable 
accuracy,  and  it  follows  from  the  results  found  that  the 
atomic  and  molecular  weights  are  the  same;  that  is,  the 
molecule  of  each  in  the  state  of  vapor  contains  but  one 
atom. 

ZINC. 

Occurrence.  The  metal  occurs  in  many  parts  of  the 
world  as  blende,  ZnS,  calamine,  ZnCO3,  siliceous  calamine 


GENERAL  CHEMISTRY.  329 

Zn2SiO4.H2O,  franklinite,  an  oxide  of  zinc,  iron  and 
manganese,  and  in  other  ores.  The  first  is  the  most 
important. 

History.  Zinc  was  not  known  in  pure  condition  to 
the  ancients,  although  brass,  its  alloy  with  copper,  was. 
This  was  made  by  smelting  copper  with  certain  ores,  the 
composition  of  which  was  not  understood.  In  the  sixteenth 
and  seventeenth  centuries  several  writers  mention  zinc,  but 
how  it  was  obtained  or  from  what  ores  is  not  now  apparent. 
It  seems  to  be  certain  that  the  metal  was  first  separated  in 
quantity  in  England  about  1730,  but  the  process  was  kept 
a  secret  for  many  years.  In  3799  works  were  established 
in  Germany  and  later  elsewhere. 

Metallurgy.  The  metal  is  always  reduced  from  the 
oxide,  the  common  sulphide  being  first  brought  into  that 
condition  by  roasting  in  a  current  of  air.  The  oxide  is 
packed  into  small  retorts  with  fine  coal  and  subjected  to 
distillation,  a  large  number  of  these  retorts  being  heated 
at  one  time  in  a  furnace.  At  a  high  heat  this  reaction 
follows  : 


The  temperature  required  for  this  reaction  is  so  high 
that  the  liberated  metal  distills  from  the  retorts  into  small 
receivers  attached,  where  it  condenses. 

Zinc  is  produced  mainly  in  the  United  States,  France, 
Belgium,  Germany  and  Great  Britain.  The  world's  pro- 
duction for  1896  was  463,444  tons,  of  which  this  country 
furnished  77,475  tons. 

Properties.  Zinc  is  a  light  colored  metal  with  a 
peculiar  bluish  tinge.  In  thick  pieces  it  is  brittle  at  the 
ordinary  temperature.  Above  100°  it  becomes  readily 
malleable,  and  at  130°  it  may  be  rolled  out  into  sheets.  At 
200°  it  becomes  brittle  again  and  may  be  powdered  in  a 
mortar.  The  melting  and  boiling  temperatures  are  given 
in  the  table  above.  It  is  acted  on  but  slightly  by  dry  or 
moist  air,  but  dissolves  easily  in  acids,  and  also  in  strong 
alkali  solutions.  With  the  last  it  forms  zincates. 


330  ,.  GENERAL  CHEMISTRY. 

Uses.  Zinc  is  largely  used  in  sheet  form  for  roofing 
and  manufacture  of  many  articles.  As  a  coating  on  iron  it 
forms  galvanized  iron.  With  copper  it  yields  several 
alloys,  brass  being  the  best  known.  In  galvanic  batteries 
it  is  usually  the  active  element,  and  lor  this  purpose  large 
quantities  are  consumed.  It  is  employed  in  the  precipita- 
tion of  certain  metals,  less  basic  in  their  character,  as  illus- 
trated by  this  experiment : 

Ex.  166.  Pour  a  dilute  solution  of  copper  sulphate  into  a  beaker, 
and  add  to  it  some  small  fragments  of  zinc.  In  a  short  time  the  zinc 
becomes  covered  with  a  spongy  coating  of  copper.  If  time  enough  be 
given  all  of  the  copper  will  disappear  from  the  solution,  as  shown  by 
loss  of  color,  a  corresponding  amount  of  zinc  taking  its  place,  as 
indicated  by  this  reaction  : 

Zn-|~CuSO4=ZnSO4-f-Cu. 
A  somewhat  similar  behavior  with  lead  will  be  later  illustrated. 

Zinc  Oxide.  This  substance  is  known  as  zinc  white 
and  is  made  in  large  quantities  by  direct  oxidation  of  zinc 
in  a  current  of  air.  Used  as  a  pigment  because  of  its  per- 
manent white  color  even  in  atmospheres  containing  traces 
of  hydrogen  sulphide.  When  heated  the  oxide  becomes 
yellow,  but  on  cooling  it  returns  to  white. 

Zinc  Hydroxide  is  obtained  by  the  precipitation  of  a 
soluble  zinc  salt  by  an  alkali  hydroxide.  In  an  excess  of 
the  reagent  the  precipitate  dissolves,  forming  a  zincate,  as 
K2ZnO2. 

Zinc  Chloride,  ZnCl2,  is  an  important  salt  formed  by 
the  solution  of  the  metal  in  hydrochloric  acid,  also  in  an- 
hydrous condition  by  distilling  a  mixture  of  ZnSO4  with 
CaCl2.  The  solution  is  used  largely  in  the  impregnation 
of  wood  for  preservation  and  also  as  a  deodorizer  and  dis- 
infectant. The  anhydrous  salt  is  employed  in  organic 
experimental  chemistry  because  of  its  marked  power  of 
absorbing  water  and  thus  bringing  about  certain  combina- 
tions. Zinc  chloride  can  be  volatilized  under  conditions 
which  permit  a  determination  of  its  vapor  density. 

Zinc  Sulphate.       In    crystallized    form    this    salt    is 


GENERAL  CHEMISTRY.  331 

known  as  white  vitriol,  ZnSO4.7H2O.  It  is  easily  made 
by  dissolving  zinc  in  dilute  sulphuric  acid,  and  is  very 
soluble  in  water.  Many  other  zinc  preparations  are  made 
from  this  salt.  It  is  isomorphous  with  Epsom  salt,  MgSO4. 
7H2O. 

Among  other  zinc  compounds  the  sulphide,  ZnS, 
and  the  carbonate,  ZnCO3,  may  be  mentioned.  The  first 
is  a  white  precipitate  obtained  by  action  of  ammonium  sul- 
phide on  a  soluble  zinc  salt.  It  is  readily  soluble  in  weak 
mineral  acids.  The  second  is  obtained  in  impure  condi- 
tion by  precipitating  soluble  zinc  salts  by  alkali  car- 
bonates. 

Recognition.  The  precipitation  of  the  sulphide  and 
carbonate,  just  referred  to,  may  be  employed  in  the  test- 
ing for  zinc.  The  formation  of  a  white  hydroxide,  which 
dissolves  in  excess  of  the  precipitant,  is  also  valuable  as  a 
test. 

CADMIUM. 

Occurrence.  This  metal  is  usually  found  as  sulphide 
with  zinc  sulphide,  and  also  in  other  zinc  ores,  but  not  in 
large  amount.  A  pure  sulphide,  CdS,  occurs  as  a  rare 
mineral. 

History  and  Preparation.  The  metal  was  discovered 
in  1818  as  a  constituent  of  crude  zinc.  As  it  is  much  more 
volatile  than  zinc  it  separates  and  distills  first  in  the  usual 
process  of  obtaining  that  metal,  explained  above.  By  re- 
peated distillation  the  metal  may  be  obtained  practically 
free  from  zinc. 

Properties  and  Uses.  It  is  tough  and  somewhat 
malleable,  standing  between  zinc  and  lead  in  hardness.  It 
forms  readily  a  number  of  alloys,  some  of  which  have  a 
very  low  melting  point. 

The  salts  of  cadmium  are  not  important.  The  sulphide, 
CdS,  is  a  yellow  pigment;  the  iodide,  CdI2,  is  a  heavy  salt 
sometimes  used  in  making  photographic  emulsions.  The 
chloride,  CdCl2,  and  the  sulphate,  CdSO4,  are  easily  made 
by  solution  of  the  metal  in  hydrochloric  or  sulphuric  acid. 


332  GENERAL  CHEMISTRY. 

Recognition.  The  salts  of  cadmium  yield  a  yellow 
precipitate  with  hydrogen  sulphide  or  ammonium  sulphide 
which  is  not  soluble  in  an  excess  of  alkali  sulphide.  Cad- 
mium hydroxide,  CdO2H2,  is  a  white  precipitate  obtained 
by  adding  an  alkali  hydroxide  to  a  cadmium  solution.  It 
is  soluble  in  an  excess  of  ammonia  water. 

MERCURY. 

Occurrence.  The  native  metal  is  found  in  small  quan- 
tity, but  the  important  ore  is  the  red  sulphide  or  cinnabar, 
HgS.  This  is  mined  in  California,  in  Spain,  in  the  Aus- 
trian province  of  Carniola  and  elsewhere.  The  world's 
production  in  1896  was  4,080  tons,  of  which  the  United 
States  furnished  1,300  tons  and  Spain  1,633  tons. 

History.  Mercury  was  known  to  the  Greeks  and  early 
Latin  writers;  but  its  importance  was  not  great  until  the 
time  of  the  alchemists,  who  studied  its  properties.  Because 
of  its  numerous  uses  in  physical  and  chemical  researches, 
its  properties  have  been  carefully  investigated  in  the  pres- 
ent century  by  many  scientists. 

Metallurgy.  The  metal  is  usually  separated  by  dis- 
tillation of  the  sulphide  with  lime  in  iron  retorts,  or  by  the 
action  of  hot  gases  from  the  combustion  of  fuel  on  the  ore 
in  furnaces  of  peculiar  construction. 

In  this  case  the  sulphur  burns  to  SO2  and  the  mercury 
is  liberated,  to  be  condensed  in  long,  tight  chambers  con- 
nected with  the  furnace. 

Properties.  This  is  the  only  metal  liquid  at  the  or- 
dinary temperature.  It  boils  at  357°  C.  and  at  — 39°  be- 
comes solid.  The  rate  of  expansion  by  heat  is  very  regu- 
lar for  mean  ranges  of  temperature,  being  0.0001815  of  the 
volume  at  0°  for  each  degree.  For  higher  temperatures 
the  rate  of  expansion  increases  but  slightly,  upon  which 
fact  the  value  of  the  metal  in  constructing  accurate  ther- 
mometers partly  depends.  As  it  may  be  easily  distilled 
the  purification  is  not  difficult,  and  the  metal  may  therefore 
be  brought  into  a  condition  of  constant  composition  es- 


GENERAL  CHEMISTRY.  333 

sential  in  any  substance  employed  as  a  standard.  The 
important  unit  of  electrical  resistance  is  based  on  the  re- 
sistance of  a  column  of  pure  mercury  of  certain  dimensions. 
The  metal  unites  readily  with  many  other  metals,  form- 
ing mixtures  called  amalgams. 

Uses.  Mercury  is  used  largely  in  the  extraction  of 
gold  and  silver  by  amalgamation,  already  referred  to.  It 
is  employed  in  the  construction  of  thermometers,  barom- 
eters, pressure  gauges  and  many  other  instruments  for  sci- 
entific and  technical  measurements.  The  amalgams  are 
very  useful;  one  with  tin  serves  for  the  backing  of  mirrors, 
some  with  cadmium  and  tin  and  also  other  metals  are  used 
in  filling  teeth,  while  other  combinations  find  different 
applications  in  the  arts. 

Compounds  of  Mercury.  Like  copper,  mercury 
forms  two  series  of  compounds,  the  mercurous  and  mer- 
curic, some  of  which  find  extensive  use  in  medicine. 

Mercury  Oxides.  Mercurous  oxide,  Hg2O,  is  a  black 
powder  formed  by  precipitation  of  a  mercurous  salt  by 
solution  of  KOH,  or  by  digesting  calomel,  Hg2Cl2,  with 
this  alkali.  It  oxidizes  readily,  and  on  standing  decom- 
poses partly  into  mercury  and  mercuric  oxide.  This  oxide, 
HgO,  is  obtained  by  heating  mercury  in  the  air  to  about 
300°  to  350°,  also  by  calcining  the  nitrate.  A  crystalline 
red  powder  is  made  in  these  processes,  and  this  has  been 
used  in  experiments,  already  given,  on  the  liberation  of 
oxygen.  By  precipitating  mercuric  chloride  with  potassium 
hydroxide  an  amorphous  yellow  oxide  is  obtained. 

Mercuric  Sulphide,  HgS,  is  obtained  as  a  black  pre- 
cipitate by  passing  H2S  into  a  solution  of  a  mercuric  salt. 
When  this  precipitate  is  dried  and  sublimed  it  turns  red. 
A  red  sulphide,  called  vermilion,  is  made  directly  by  sub- 
liming a  mixture  of  mercury  or  mercuric  oxide  with  sul- 
phur. It  is  used  as  a  valuable  red  pigment,  and  is  identi- 
cal in  composition  with  the  native  cinnabar.  Mercurous 
sulphide  cannot  be  obtained  by  precipitation  of  mercurous 
compounds  by  hydrogen  sulphide. 


334  GENERAL  CHEMISTRY. 

Mercury  Chlorides.  Mercurous  chloride,  Hg2Cl2, 
is  commonly  known  as  calomel  and  is  made  by  subliming 
an  intimate  mixture  of  mercury  and  corrosive  sublimate 
which  have  been  well  rubbed  together.  It  is  an  amorphous 
white  powder,  insoluble  in  water,  and  is  largely  used  in 
medicine,  the  sublimed  mass  being  prepared  for  this  pur- 
pose by  grinding  arid  thorough  washing  with  water  to 
remove  any  of  the  other  chloride  carried  over.  Mercurous 
chloride  turns  black  by  the  addition  of  ammonia  water. 
Mercuric  chloride,  HgCl2,  is  known  as  corrosive  sublimate 
and  is  usually  made  by  distilling  a  mixture  of  common 
salt  and  mercuric  sulphate.  It  is  a  white,  crystalline  mass, 
soluble  in  water,  and  violently  poisonous.  With  ammonia 
it  yields  a  white  precipitate: 

HgCl2+2NH3:=NH4Cl+HgClNH2. 

This  chloride  is  largely  used  in  the  sterilization  of  sur- 
gical instruments  before  and  after  use,  and  in  the  disinfec- 
tion of  articles  which  have  been  exposed  to  contagion.  A 
solution  of  1  part  to  1,000  or  2,000  of  water  is  sufficient  for 
such  purposes.  Steel  instruments  are,  however,  corroded 
by  the  solution.  It  is  readily  soluble  in  glycerol,  and  this 
solution  is  often  used  in  the  preparation  of  aqueous  solu- 
tions by  dilution. 

Mercuric  Iodides.  Mercurous  iodide,  Hg2I2,  may  be 
obtained  as  a  green  mass  by  rubbing  mercury  and  iodine 
together  in  a  mortar  in  proper  proportions.  It  is  also  made 
by  precipitation  of  mercurous  nitrate  by  a  solution  of  po- 
tassium iodide.  It  is  not  very  stable.  Mercuric  iodide, 
HgI2,  is  a  beautiful  bright  red  precipitate  obtained  by  add- 
ing potassium  iodide  solution  in  right  amount  to  a  mer- 
curic solution: 

2KI+HgCl8=2KCl+HgI8. 

The  precipitate  is  readily  soluble  in  an  excess  of  the  po- 
tassium iodide,  and  the  solution  so  obtained  is  a  valuable 
laboratory  reagent.  The  iodide  may  be  sublimed  at  a  high 
temperature. 

Other   Compounds.     Mercuric  sulphate,   HgSO4,    is 


GENERAL  CHEMISTRY.  335 

made  by  solution  of  mercury  in  strong  sulphuric  acid.  It 
serves  for  the  preparation  of  the  chloride.  By  treating 
with  boiling  water  it  is  converted  into  a  yellow  basic  salt 
known  as  turpeth  mineral,  HgSO4. 2HgO.  Mercuric  nitrate, 
Hg(NO3)2,  is  made  by  action  of  strong  nitric  acid  on 
mercury;  with  a  weak  acid  in  the  cold  mercurous  nitrate 
Hg2(NO3)2  is  obtained. 

Recognition.  Mercury  in  compounds  may  be  recog- 
nized by  the  separation  of  the  metal  as  a  bright  amalgam 
when  a  copper  wire  or  bit  of  foil  is  rubbed  with  some  of 
the  substance  mixed  with  hydrochloric  acid.  In  solid 
bodies  it  may  be  easily  detected  by  fusion  with  a  mixture 
of  dry  sodium  carbonate  and  saltpeter  in  a  glass  tube.  The 
volatile  metal  is  liberated  and  condenses  in  droplets  on  the 
cooler  upper  part  of  the  tube. 


CHAPTER  XVII. 


BORON,  ALUMINUM,  GALLIUM,  INDIUM,  THALLIUM, 

SCANDIUM,  YTTRIUM,  LANTHANUM 

AND  YTTERBIUM. 

GENERAL   CHARACTERISTICS. 

THESE  elements  form  the  third  family  in  the  Periodic 
System  and  with  three  exceptions  are  technically  quite 
unimportant.  The  compounds  of  boron  have  been  referred 
to  already,  while  those  of  aluminum  will  be  briefly  de- 
scribed below.  A  few  thallium  compounds  are  somewhat 
important,  while  those  of  the  other  elements  are  very 
rare.  The  chief  physical  properties  are  shown  in  the  fol- 
lowing table: 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

Boron.          

10.95 

2.68 

4.1 

high 

Aluminum    

27.11 

2.60 

10.4 

700° 

Gallium       

69.91 

5.95 

11.7 

30° 

Indium  

113.85 

7.42 

15.3 

176° 

Thallium  

204.15 

11.86 

17.2 

290° 

Scandium 

44  12 

Yttrium                         ...    . 

8902 

Lanthanum           

138.64 

6.20 

22  3 

Ytterbium  

173.19 

The  nine  elements  in  this  family  fall  naturally  into  two 
groups;  a  primary  group  containing  scandium,  yttrium, 
lanthanum  and  ytterbium,  and  a  secondary  group  contain- 
ing boron,  aluminum,  gallium,  indium  and  thallium.  With 


GENERAL  CHEMISTRY.  337 

the  exception  of  boron  they  behave  in  their  best  known 
compounds  as  metals  and  are  all  trivalent.  Thallium  is 
closely  related  to  the  alkali  metals  on  the  one  hand  and  to 
lead  on  the  other,  but  it  is  usually  placed  in  this  group.  It 
will  be  recalled  that  the  metallic  character  of  the  elements 
in  the  horizontal  periods  diminishes  on  passing  from  left  to 
right  in  the  table  of  the  Periodic  arrangement;  we  have  in 
boron  the  first  of  the  elements  in  which  the  acid  behavior 
is  quite  marked.  In  the  next  family  the  change  will  be 
found  still  more  characteristic. 


ALUMINUM. 

Occurrence.  This  metal  is  very  abundant  in  combina- 
tion. It  is  found  as  silicate  in  all  clay  and  in  more  or  less 
pure  condition  in  several  other  minerals,  as  the  felspars, 
AlKSi3O8  and  AlNaSi3O8.  It  occurs  as  oxide,  A12O3, 
in  corundum,  ruby,  sapphire  and  emery;  as  hydroxide  or 
basic  hydroxide  in  several  minerals  of  which  bauxite, 
A12O(OH)4,  is  the  most  important;  as  fluoride  in  cryolite, 
AlF3.3NaF.  Clays,  which  are  crude,  hydrated  silicates, 
are  formed  by  the  weathering  or  disintegration  of  felspars 
and  similar  silicates.  . 

History.  The  important  properties  of  the  alums  were 
known  to  the  alchemists,  and  over  one  hundred  years  ago 
it  was  recognized  that  a  peculiar  earth  is  combined  with 
sulphuric  acid  in  these  bodies.  In  his  important  elec- 
trical decompositions  Davy,  early  in  this  century,  tried  to 
separate  the  metal  which  he  believed  this  earth  must  con- 
tain, but  without  success.  In  1827  Woehler  succeeded  in 
isolating  the  metal  by  decomposing  the  chloride  by  means 
of  potassium: 

A1C13  +  3K  =  3KC1+A1. 


Metallurgy.  For  many  years  all  the  aluminum  of 
commerce  was  made  by  a  process  based  on  the  above  reac- 
tion. Potassium  was  replaced  by  sodium  and  great 
improvements  were  made  in  the  methods  of  producing  the 
chloride,  A1C13,  but  at  best  the  processes  were  expensive 


338  GENERAL  CHEMISTRY. 

and  the  metal  found  but  limited  applications.  Recently, 
methods  have  been  perfected  by  which  the  reduction  is 
effected  through  the  agency  of  the  electric  current.  The 
electric  furnace  in  which  this  reduction  is  carried  out  is  a 
large  iron  box  lined  with  carbon.  In  this  is  placed  a 
mixture  of  powdered  cryolite  and  carbon  and  a  powerful 
current  is  passed  in  through  massive  anodes  of  hard  car- 
bon, the  box  itself  serving  as  the  cathode.  The  cryolite 
is  fused  by  the  intense  heat  generated  by  the  passage  of 
the  current,  and  then  powdered  alumina,  made  from 
bauxite,  is  added  from  time  to  time.  This  is  decomposed 
into  metal  and  oxygen,  the  former  sinking  to  the  bottom 
of  the  box  or  furnace,  from  which  it  is  drawn  off  from  time 
to  time,  while  the  latter  is  taken  up  by  the  carbon  of  the 
bath  and  of  the  anodes.  The  reaction  may  be  explained 
as  one  of  electrolysis,  or,  on  the  other  hand,  it  may  be  con- 
sidered as  one  between  carbon  and  alumina,  A12O3,  made 
possible  by  the  high  heat  between  the  electrodes.  In 
some  of  the  later  modifications  of  the  electric  furnace 
method  aluminum  of  a  high  degree  of  purity  is  produced. 
In  1896  the  United  States  produced  1,300,000  pounds 
of  aluminum,  France  about  600,000  pounds  and  Switzer- 
land about  1,450,000  pounds. 

Properties.  Aluminum  is  a  white,  malleable  and 
ductile  metal,  comparatively  soft  and  not  easily  corroded 
in  the  air  or  by  acid  fumes.  Even  at  a  high  temperature 
it  oxidizes  but  imperfectly.  At  a  red  heat  it  decomposes 
steam.  Nitric  and  sulphuric  acids  attack  it  but  slowly, 
but  in  hydrochloric  acid  it  dissolves  readily.  With  potas- 
sium or  sodium  hydroxide  it  forms  analuminate,  with  liber- 
ation of  hydrogen.  Aluminum  alloys  with  a  number  of 
metals;  with  copper  it  forms  a  valuable  bronze  called 
aluminum  bronze. 

Uses.  Because  of  its  low  specific  gravity  and  resist- 
ance to  oxidation  the  metal  has  come  into  use  for  the  man- 
ufacture of  numerous  small  articles,  such  as  beams  of 
balances,  handles  of  surgical  instruments,  combs  and 
backs  of  brushes,  buckles  and  even  cooking  utensils.  It 


GENERAL  CHEMISTRY.  339 

was  at  one  time  predicted  that  the  metal  would  come  into 
use  for  structural  purposes,  in  place  of  iron,  but  these  ex- 
pectations have  not  been  realized. 

Aluminum  Oxide.  This  is  found  in  nature  in  several 
forms,  as  mentioned  above.  The  ruby  and  the  sapphire 
have  also  been  formed  artificially.  The  cheaper  emery  is 
commonly  used  for  grinding  and  polishing  metals  and  glass, 
as  it  is  extremely  hard.  Amorphous  aluminum  oxide  may 
be  easily  made  by  precipitating  the  hydroxide,  A1(OH)3, 
washing  it  thoroughly  and  finally  heating  to  a  high  tem- 
perature. It  is  left  in  this  way  as  a  white  substance  quite 
insoluble  in  water,  and  but  slowly  soluble  in  acids. 

Aluminum  Hydroxides.  The  simple,  normal  hydrox- 
ide, A1(OH)3,  is  made  as  a  laboratory  product: 

=  2A1(OII)34-3(NH4)2SO4. 


As  so  made  it  is  a  white,  gelatinous  mass.  In  nature 
several  hydroxides  occur,  the  most  important  being  baux- 
ite. In  this,  part  of  the  aluminum  is  usually  replaced  by 
iron  giving  bodies  varying  between  A12O(OH)4  and 
AlFeO^OH)4.  The  varieties  low  in  iron  are  used  in  the 
manufacture  of  the  alums  and  the  metal  itself. 

Aluminum  Chloride.  This  is  easily  obtained  in  solu- 
tion by  dissolving  the  metal  or  hydroxide  in  hydrochloric 
acid.  The  pure,  dry  substance  is  made  by  passing  dry 
chlorine  gas  over  a  heated  mixture  of  aluminum  oxide  and 
charcoal,  or,  better,  by  passing  dry  chlorine  over  hot  alumi- 
num turnings,  in  a  tube  heated  in  a  gas  furnace.  When 
obtained  in  this  way  it  is  a  solid,  colorless  substance  which 
may  be  vaporized  at  a  moderately  high  temperature.  The 
density  of  the  vapor  corresponds  to  the  molecule,  A12C16; 
but  at  a  higher  temperature  the  density  is  such  as  to  giye 
A1C13  as  the  formula,  suggesting  the  trivalent  character  of 
the  metal.  The  chloride  is  a  valuable  agent  in  organic 
synthesis. 

Aluminum  Sulphate.  As  the  group  SO4  is  bivalent 
and  aluminum  trivalent,  the  formula  of  this  salt  is 


340  GENERAL  CHEMISTRY. 

A12(SO4)3.  It  is  usually  made  by  dissolving  the  hydrox- 
ide from  bauxite  in  dilute  sulphuric  acid.  The  evaporated 
liquid  yields  the  salt  with  a  large  amount  of  combined 
water.  This  product,  or  one  more  completely  evaporated, 
is  used  for  several  purposes  in  the  arts,  especially  in  the 
clarification  of  water,  as  a  mordant  in  dyeing  and  in  the 
glazing  of  paper.  It  is  also  used  in  the  manufacture  of 
some  of  the  common  alums. 

The  Alums.  Aluminum  sulphate  combines  with 
alkali  sulphates  and  water,  forming  bodies  of  the  following 
composition  : 

K2SO4.A12(SO4)3.24H2O 

Na2SO4.Al2(SO4)3.24H2O 

(NH4)2S04.A12(S04)3.24H20. 

These  are  called  alums,  the  first  and  trfird  being  com- 
mon commercial  substances.  As  they  crystallize  well,  they 
may  be  easily  purified,  which  is  not  the  case  with  the  very 
soluble  aluminum  sulphate.  These  alums  are  often  used 
in  the  place  of  the  simple  sulphate,  and  may  be  made  from 
several  other  substances  than  bauxite  as  a  starting 
material. 

The  term  alum  is  applied  to  a  large  number  of  double 
sulphates  besides  those  containing  aluminum.  Thus, 
chrome  alum  is  K2SO4.Cr2(SO4)3.24H2O;  ferric  alum  is 
(NH4)2SO4.Fe2(SO4)3.24H2O.  In  general,  an  alum  is 
a  double  sulphate,  one  part  being  an  alkali  sulphate 
(Li,  Na,  K,  Cs,  Rb,  NH4),  and  the  other  part  a  sulphate 
of  a  trivalent  metal  (usually  Al,  Fe  or  Cr)  with  24H2O. 
These  alums  are  isomorphous  and  crystallize  in  octahedra. 
It  is  a  further  interesting  fact  that  the  analogous  selenates 
and  tellurates  belong  to  the  same  isomorphous  group. 

Aluminum  Silicates.  These  occur  in  nature  in  a  great 
number  of  minerals.  Kaolin  is  nearly  pure  aluminum 
silicate,  while  clay  is  a  mixture  in  which  this  silicate  largely 
predominates.  These  bodies  are  used  in  the  production 
of  articles  of  earthenware  or  pottery  in  various  forms. 
The  finer  pottery  known  as  porcelain  is  made  of  the  purest 
kaolin,  while  other  kinds  are  made  of  commoner  material. 


GENERAL  CHEMISTRY.  341 

Pottery. 

In  the  manufacture  of  pottery  the  clay  or  kaolin  to  be 
used  is  ground  to  a  fine  powder  which  is  mixed  with  some 
powdered  material  to  serve  as  a  flux.  This  is  often  felspar, 
or  it  may  be  a  mixture  of  chalk  or  gypsum  with  quartz. 
Pure  clay  does  not  fuse,  even  at  a  very  high  temperature, 
and  the  flux  is  employed  to  melt,  surround  the  clay 
particles  and  bind  them  together.  To  this  end  the  pre- 
pared mixed  material  is  ground  up  with  water  to  form  a 
smooth,  homogeneous  mass,  which  is  molded  into  the 
desired  articles.  These  are  dried  and  gradually  heated  to  a 
high  temperature,  by  which  a  more  or  less  porous  product 
known  as  biscuit  ware  is  obtained.  For  many  purposes 
this  ware  cannot  be  employed,  but  must  be  covered  with  a 
glaze.  In  some  cases,  to  secure  this,  the  articles  are 
covered  with  a  paste  containing  litharge  and  reheated,  by 
which  a  smooth  lead  silicate  is  formed  and  flows  over  the 
whole  mass.  A  so-called  salt  glaze  is  produced  on  some 
kinds  of  cheaper  wares  by  throwing  salt  into  the  kilns  in 
which  the  biscuit  or  open  ware  is  fired,  just  before  the 
end  of  the  operation.  This  volatilizes  and  is  decomposed, 
yielding  finally  a  fusible  sodium-aluminum  silicate  to  flow 
over  the  whole  surface.  Mixtures  of  quartz,  chalk  and 
borax  are  used  in  some  glazes.  For  hard  laboratory  por- 
celain a  mixture  of  kaolin,  quartz,  lime  and  broken  por- 
celain is  often  used.  When  this  is  fused  over  the  fine 
biscuit  ware  a  smooth,  hard  surface  is  secured  which  is 
intended  to  resist  the  action  of  chemicals. 

Other  Compounds.  Cryolite  is  a  double  fluoride  hav- 
ing the  composition  AlF3.3NaF.  It  is  used  in  the  metal- 
lurgy of  aluminum,  also  to  produce  sodium  carbonate  and 
for  its  fluxing  properties  in  general.  Ultramarin  is  a  com- 
plex mixture  obtained  by  fusing  kaolin  with  sodium  sul- 
phate and  carbon,  or  with  sulphur  and  sodium  carbonate. 
The  composition  is  variable,  but  one  kind  has  approxi- 
mately the  formula  Al3Si2Na2SO9.  Aluminum  acetate  is 
used  as  a  valuable  mordant  in  dyeing.  Aluminates  may 
be  formed  by  solution  of  the  metal  or  hydroxide  in  caustic 


342  GENERAL  CHEMISTRY. 

alkalies,  or  by  fusion  of  the  hydroxide  with  caustic  or  car- 
bonate alkali. 

Ex.  167.  Place  some  bits  of  aluminum  wire  in  a  test-tube  and  add 
solution  of  sodium  hydroxide.  A  slow  evolution  of  hydrogen  gas  begins, 
which  is  hastened  by  application  of  heat.  The  metal  dissolves  to  form 
sodium  aluminate,  NaAlO2. 

Recognition.  The  aluminum  in  soluble  compounds  is 
easily  recognized  by  the  white  precipitate  of  hydroxide 
formed  on  the  addition  of  ammonia  water.  This  precip- 
itate is  soluble  in  an  excess  of  caustic  soda  or  potassa 
solution. 

GALLIUM  AND  INDIUM. 

These  are  two  very  rare  metals  found  in  the  zinc  blendes 
of  certain  localities.  Both  were  discovered  by  aid  of  the 
spectroscope,  the  former  in  1875  by  L.  de  Boisbaudran, 
and  the  latter  by  Reich  and  Richter  in  1863.  Gallium 
melts  at  30°  C.,  that  is  below  the  body  temperature,  and  is 
characterized  by  the  two  blue  violet  lines  it  shows  in  the 
spectrum.  Indium  shows  likewise  two  blue  lines  in  the 
spectrum.  Both  metals  behave  as  trivalent  in  most  of 
their  compounds,  and  both  form  alums,  and  several  other 
compounds  resembling  those  of  aluminum. 

THALLIUM. 

Occurrence  and  History.  This  comparatively  rare 
metal  is  found  in  certain  iron  and  copper  pyrites  and  in 
a  few  other  minerals.  It  was  discovered  in  1861  by 
Crookes  in  the  examination  of  a  deposit  from  the  lead 
chambers  of  a  German  sulphuric  acid  factory  where  pyrites 
were  burned  to  furnish  the  sulphurous  oxide.  It  is  pre- 
pared usually  from  similar  deposits. 

Properties.  Thallium  is  a  soft  metal  resembling  lead 
in  appearance  and  in  physical  properties.  But  it  oxidizes 
in  moist  air  yielding  an  oxide,  T12O,  which  dissolves  in 
water  with  formation  of  an  alkaline  hydroxide,  T1OH.  In 
this  behavior  it  resembles  the  alkali  metals.  It  decom- 
poses water  at  an  elevated  temperature  and  dissolves 
readily  in  dilute  sulphuric  acid. 


GENERAL  CHEMISTRY.  343 

Uses.  The  metal  has  no  technical  applications,  but 
several  of  its  compounds  are  practically  useful.  Some  of 
them  are  used  in  making  a  very  dense  glass  with  high 
refractive  index,  for  optical  purposes. 

Compounds  of  Thallium.  Some  of  these  resemble 
the  corresponding  lead  compounds  closely,  while  others 
are  similar  to  the  compounds  of  the  alkali  metals.  In  a  few 
salts  a  resemblance  to  the  aluminum  salts  is  found.  From 
these  conditions  it  is  evident  that  the  proper  place  of  this 
metal  among  the  elements  is  hard  to  define.  Two  classes 
of  compounds  are  known,  some  being  univalent,  the  others 
trivalent.  T1C1,  TIBr  and  Til  are  insoluble  bodies 
resembling  the  lead  and  silver  salts.  T1OH  and  T12SO4 
are  soluble  and  in  some  respects  similar  to  KOH  and 
K2SO4.  The  sulphate  forms  an  alum  with  A12(SO4)3.  On 
the  other  hand,  we  have  thallic  sulphate,  T12(SO4)3,  which 
combines  with  potassium  sulphate  to  form  an  alum-like 
body.  Thallous  hydroxide,  T1OH,  is  colored  brown  by 
ozone,  and  paper  moistened  in  such  a  solution  is  some- 
times used  as  an  ozone  test. 

Recognition.  Thallium  compounds  are  easily  recog- 
nized by  the  bright  green  line  they  exhibit  in  spectroscopic 
examination. 

THE  RARE  EARTHS. 

Several  complex  minerals  are  known  which  contain  a 
number  of  closely  related  metals,  the  oxides  of  which  are 
spoken  of  as  the  rare  earths.  One  of  these  minerals 
is  called  gadolinite;  it  is  a  silicate  of  yttrium  and  a  dozen 
other  metals  in  smaller  proportion.  Another  mineral  is 
known  as  euxenite,  which  contains  scandium,  ytterbium 
and  several  other  metals  in  combination  with  titanic  acid. 
Lanthanum  is  found  in  these  minerals  and  in  several 
others.  The  separation  of  all  the  metals  contained  in  these 
complex  compounds  is  a  matter  of  extreme  difficulty,  and 
the  problem  cannot  as  yet  be  considered  as  completely 
solved.  Something  more  will  be  said  about  these  bodies 
in  the  following  chapter,  where  some  important  technical 
applications  of  several  related  earths  will  be  explained. 


CHAPTER  XVIII. 


THE  CARBON  GROUP :     CARBON,  SILICON,  GERMA- 
NIUM, TIN  AND  LEAD.— THE  TITANIUM  GROUP: 
TITANIUM,  ZIRCONIUM,    CERIUM  AND 
THORIUM. 

GENERAL  CHARACTERISTICS. 

THE  fourth  family  of  the  Periodic  System  contains  a 
number  of  elements  of  which  the  nine  above  given 
have  been  pretty  thoroughly  studied.  Carbon  and  silicon 
have  been  already  described,  as  fully  as  is  necessary  for 
our  purpose,  in  the  eighth  and  tenth  chapters.  In  the 
present  chapter  they  are  introduced  again  merely  to  show 
their  relations  to  the  other  members  of  the  family.  It  is 
customary  to  divide  this  family  into  two  groups,  but  the 
division  is  somewhat  arbitrary.  In  the  following  table  the 
important  physical  constants  of  the  elements  are  given  : 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

Carbon        .  .        

12  01 

8  50 

3  4 

very  high 

Silicon  

28.40 

2  49 

11  4 

very  high 

Germanium  

72.48 

5.47 

13  2 

900°  (?) 

Tin   

119.05 

7  29 

16.3 

238° 

Lead 

206  92 

11  40 

18  1 

326° 

Titanium 

48  15 

3  70 

13  0 

very  high 

Zirconium                  

90  40 

4  15 

21  7 

1500°  (?) 

Cerium      

140.20 

6.70 

20  9 

800°  (  ?) 

Thorium  

232.63 

11.20 

20  8 

very  high 

The  family  likeness  of  these  elements  may  be  illustrated 


GENERAL  CHEMISTRY.  345 

by  collecting  the  formulas  oi   some  of  their  best  known 
compounds. 

CO2  CS2  CH4  CC14 

Si02  SiS8  SiH4          SiCl4  Si(OH)4 

GeO2  GeSa  GeCl4 

SnO2  SnS2  SnCl4 

Pb02  PbCl4 

Ti02  TiS2                                TiCl4  Ti(OH)4 

ZrO2                                                     ZrCl4  Zr(OH)4 

Ce02 

Th02  ThS2                               ThCl4  Th(OH)4 

It  will  be  noticed  that  in  the  groups  of  oxides,  sul- 
phides and  chlorides  there  is  an  almost  complete  paral- 
lelism. Bromides  and  fluorides  of  the  type  MX4  are 
known  for  most  of  the  elements  in  the  list. 

The  acid  character  of  some  of  these  elements  is  very 
pronounced,  as  seen  in  the  carbonates  and  silicates. 
Stannates  and  titanates  of  the  light  metals  are  known; 
with  germanium,  lead  and  thorium  the  metallic  character 
predominates. 

GERMANIUM. 

This  is  a  very  rare  element  which  was  discovered  in 
1886  in  a  silver  ore.  It  is  sometimes  observed  in  tin  ores 
but  only  in  small  amount.  Like  tin,  it  forms  two  series  of 
salts;  in  the  first  it  is  bivalent  and  in  the  other  quadri- 
valent. 

TIN. 

Occurrence.  This  important  metal  has  been  found  in 
several  ores,  but  the  valuable  ore  is  tinstone  or  cassiterite, 
SnO2,  which  occurs  in  Cornwall,  in  Australia,  in  the 
Island  of  Banca,  in  the  Malay  Peninsula,  in  Bolivia  and 
elsewhere.  A  tin  pyrite  is  also  known  but  it  is  not  abun- 
dant. 

History.  Tin  was  known  in  very  early  times,  and  it 
appears  that  before  the  beginning  of  the  Christian  era  the 


346  GENERAL  CHEMISTRY. 

mines  or  deposits  of  Cornwall  were  worked  by  the  peoples 
living  around  the  eastern  shores  of  the  Mediterranean.  It 
was  largely  used  in  making  bronze  and  in  soldering  lead. 
The  properties  and  uses  of  tin  were  well  known  to  the 
alchemists,  and  possibly  even  before  their  time  its  value  as 
a  coating  for  iron  was  known.  This  knowledge  was  not 
applied  practically  until  about  1720,  however.  To-day 
this  is  the  most  important  use  to  which  tin  is  put. 

Metallurgy.  Tinstone  is  a  very  hard  ore  and  somewhat 
difficult  to  reduce.  Several  reactions  have  been  employed 
in  separating  the  metal,  but  now  it  is  commonly  reduced 
by  carbon  in  a  blast  furnace  or  in  a  kind  of  reverberatory 
furnace.  The  ore  is  stamped  and  washed  to  separate 
earthy  matter  and  roasted  to  convert  sulphides  into  oxides. 
So  prepared  it  is  ready  for  the  actual  reduction: 


The  crude  tin  which  is  run  from  the  hearth  of  the  rever- 
beratory furnace  or  from  the  blast  furnace  may  contain 
iron,  copper  and  traces  of  other  metals.  This  is  refined  by 
the  process  termed  liquation,  in  which  the  crude  blocks 
or  ingots  are  piled  up  in  a  reverberatory  furnace  with 
inclined  hearth,  the  metal  being  at  the  highest  part  of  the 
incline.  As  tin  is  readily  fusible  it  begins  to  flow  out 
from  the  mass  when  the  furnace  is  gently  and  gradually 
heated.  The  other  metals  require  greater  heat  for  melt- 
ing and  remain  behind  in  the  operation. 

The  purest  tin  comes  from  Banca.  The  world's  pro- 
duction of  tin  in  1895  was  91,693  tons,  of  which  58,690  tons 
came  from  the  Malay  Peninsula. 

Properties.  Tin  is  a  white  metal  which  melts  at 
about  233°.  It  is  malleable  and  ductile  at  a  temperature 
of  100°.  At  a  temperature  of  200°  it  becomes  so  brittle 
that  it  may  be  easily  powdered.  At  low  temperatures  also 
it  is  very  brittle.  It  alloys  easily  with  several  other 
metals,  and  resists  atmospheric  effects  in  a  marked  man- 
ner. A  bar  of  tin  when  bent  emits  a  peculiar  sound  called 
the  "  cry  of  the  tin." 


GENERAL  CHEMISTRY.  347 

Uses.  It  is  employed  mainly  in  coating  iron  and  cop- 
per. The  metal  to  be  covered  is  thoroughly  cleaned  and 
dried  and  dipped  in  a  bath  of  molten  tin.  Enough  clings 
to  form  a  good,  durable  coating.  Several  common  alloys 
contain  tin.  Pewter  is  a  mixture  of  about  4  parts  of  tin 
with  1  of  lead;  common  solder  contains  lead  and  tin  in 
equal  proportions;  bronze,  bell  metal,  speculum  metal  and 
gun  metal  are  essentially  tin  and  copper  alloys.  Tin 
amalgam  is  used  in  silvering  mirrors. 

Tin  Oxides.  Two  are  known.  The  monoxide,  SnO,' 
is  unimportant;  the  dioxide,  SnO2,  is  the  common  ore  of 
tin.  An  artificial  product  which  is  perfectly  white  is 
made  by  oxidizing  tin  with  nitric  acid  and  heating  the 
residue.  It  is  quite  insoluble  in  water.  When  this  oxide  is 
fused  with  sodium  hydroxide  a  soluble  salt  known  as 
sodium  stannate  is  produced.  It  has  the  composition 
Na2SnO3,  and  is  used  as  a  mordant  in  calico  printing. 
Hydroxides  of  tin  are  also  known. 

Tin  Sulphides.  Stannous  sulphide,  SnS,  may  be 
made  by  heating  a  mixture  of  tin  foil  and  sulphur;  also  by 
precipitation  of  a  solution  of  stannous  chloride  by  hydro- 
gen sulphide.  This  yields  a  brown  precipitate  which 
becomes  darker  on  drying.  The  disulphide,  SnS2,  is  an 
important  substance  made  usually  by  subliming  a  mixture 
of  tin  amalgam,  sulphur  and  ammonium  chloride.  The 
mercury  and  chloride  are  driven  off  in  the  sublimation  while 
the  tin  remains  in  yellow  scales  as  disulphide.  When  pre- 
pared in  this  way  the  product  is  known  as  mosaic  gold  and 
is  used  as  a  bronze  powder  for  many  purposes.  A  yellow 
precipitate  of  amorphous  SnS2  is  obtained  by  passing 
hydrogen  sulphide  into  an  acid  stannic  solution. 

Tin  Chlorides.  Stannous  chloride,  SnCl2,  is  obtained 
by  dissolving  tin  in  hydrochloric  acid.  On  evaporating 
the  solution  a  crystalline  compound,  SnCl2.2H2O,  sepa- 
rates. Under  the  name  of  tin  salt  this  is  used  as  a  mor- 
dant in  dyeing.  A  solution  of  stannous  chloride  is  a  valu- 
able reagent  in  the  laboratory.  Stannic  chloride,  SnCl4, 
is  a  fuming,  colorless  liquid  obtained  by  passing  chlorine 


348  GENERAL  CHEMISTRY. 

over  melted  tin.  It  is  readily  soluble  in  water,  yielding 
hydrates,  SnCl4.3H2O  and  SnCl4.5H2O,  with  proper 
amounts  of  water.  These  hydrates  are  obtained  in  solid 
crystalline  condition,  the  first  being  known  as  butter  of  tin. 
They  are  employed  as  mordants  to  produce  certain  colors 
in  dyeing.  When  used  for  this  purpose  they  are  most 
easily  made  by  dissolving  tin  in  aqua  regia,  evaporating 
and  crystallizing  from  water. 

Other  Compounds.  Bromides  and  iodides  of  tin 
corresponding  to  the  two  chlorides  are  known.  Nitrates  and 
sulphates  have  been  made,  but  they  are  without  importance 
and  are  not  stable. 

Recognition.  Tin  compounds  are  recognized  by  the 
hydrogen  sulphide  precipitates,  and  by  the  separation  of 
the  metal  when  zinc  is  added  to  a  tin  solution  containing 
a  little  hydrochloric  acid.  Stannous  chloride  reacts  with 
mercuric  chloride  in  this  way,  to  give  first  a  white  precip- 
itate of  calomel  and  finally  a  dark  gray  precipitate  of  me- 
tallic mercury: 


LEAD. 

Occurrence.  This  metal  occurs  usually  as  the  gray 
sulphide,  or  galena,  PbS.  It  is  found  also  as  native  lead 
in  small  amount,  and  in  rare  minerals  as  oxide,  carbonate, 
sulphate,  chromate  and  chloride.  Galena  is  found  very 
widely  distributed  and  is  especially  abundant  in  the  United 
States. 

History.  This  is  one  of  the  seven  metals  known  to 
the  peoples  of  antiquity,  and  crude  methods  were  employed 
to  reduce  it  from  the  sulphide.  In  their  northern  con- 
quests the  Romans  introduced  the  knowledge  of  smelting 
in  the  countries  where  the  ore  was  found.  The  develop- 
ment of  the  modern  methods  of  smelting  has  been  gradual. 

Metallurgy.     Galena  is    reduced    by  several   distinct 


GENERAL  CHEMISTRY.  349 

methods.  In  one  the  ore  is  heated  in  a  furnace  with  iron, 
which  takes  the  sulphur  and  leaves  the  lead  in  the  free 
state.  In  a  modification  of  this  process  the  ore  is  heated 
with  a  mixture  of  iron  oxide  and  coke  in  a  blast  furnace. 
The  iron  oxide  becomes  reduced  to  spongy  iron,  which  acts 
readily  on  the  galena.  In  a  second  method  of  smelting  the 
lead  sulphide  is  roasted  in  a  reverberatory  furnace  to  form 
a  mixture  of  oxide  and  sulphate,  leaving  a  part  of  the 
galena  unchanged.  In  a  second  stage  of  the  process  the 
heat  of  the  furnace  is  increased  and  the  excess  of  air  shut 
off;  then  the  three  substances  react  on  each  other  in  this 
way: 

2PbO+PbS  =  3Pb-fSO2 


The  metallic  lead  collects  on  the  hearth  of  the  fur- 
nace, from  which  it  is  run  off  from  time  to  time. 

At  the  present  time  in  this  country  a  large  part  of  the 
lead  is  obtained  in  the  smelting  of  ores  which  contain  gold 
and  silver  also.  This  is  done  in  a  blast  furnace  with  coke, 
using  such  a  mixture  of  ores  as  will  best  furnish  a  liquid 
slag  and  permit  a  separation  of  the  lead  and  precious 
metals  present,  as  was  explained  in  the  fourteenth  chap- 
ter. It  was  further  explained  how  the  crude  lead  bullion 
obtained  is  desilverized  and  softened  and  made  ready  for 
commerce. 

In  1896  the  world's  production  of  lead  was  about  682,- 
000  tons,  of  which  the  United  States  furnished  174,792 
tons,  Spain  187,870  tons  and  Germany  120,170  tons. 

Properties.  Lead  is  a  soft  metal,  malleable,  but  not 
very  ductile.  It  is  not  corroded  to  a  great  extent  in  the 
air  and  in  water  only  under  certain  conditions.  Ordinary 
hard  waters  have  but  little  action,  while  soft  waters  con- 
taining carbonic  acid  or  much  chloride  or  nitrate  dissolve 
it  appreciably.  Such  waters  should  not  be  conveyed 
through  lead  pipes  for  household  use.  The  metal  alloys 
readily  with  several  others.  It  is  but  slightly  attacked  by 
sulphuric  acid  and  only  slowly  by  the  other  mineral  acids. 

Solutions  of  lead  are  reduced  by  several  metals,  nota- 
bly by  zinc,  which  can  be  illustrated  as  follows: 


350  GENERAL  CHEMISTRY. 

Ex.  168.  Fill  a  bottle  with  a  very  dilute  solution  of  lead  acetate  in 
distilled  water,  and  immerse  in  it  a  thin  strip  of  sheet  zinc,  cut  and 
spread  to  represent  the  branches  of  a  tree.  Bend  the  upper  end  of  the 
zinc  so  that  it  may  hang  over  a  glass  rod  resting  on  the  neck  of  a  bottle. 
At  the  end  of  twenty-four  hours  a  crystalline  deposit  of  the  lead  in  the 
form  of  thin,  glistening  plates  appears,  suspended  from  the  branches  of 
the  zinc.  This  is  called  the  lead  tree  and  is  formed  best  when  the  solu- 
tion is  dilute  and  the  deposition,  in  consequence,  slow. 

Uses.  Lead  is  employed  in  sheet  form  for  many  pur- 
poses, especially  in  lining  the  chambers  in  which  sulphuric 
acid  is  made.  As  pipe  it  is  used  to  convey  water.  Or- 
dinary shot  consists  of  lead  frequently  alloyed  with  a  frac- 
tion of  a  per  cent  of  arsenic.  Type  metal  is  an  alloy  of 
antimony  and  lead,  with  tin  sometimes  added.  Common 
solder  is  an  alloy  of  lead  and  tin.  Several  important 
easily  fusible  alloys  contain  lead,  bismuth  and  tin. 

Lead  Oxides.     Three  are  well  known. 

Litharge,  PbO,  is  a  yellow  substance  obtained  by  oxi- 
dizing lead  in  the  air.  If  the  temperature  is  kept  below 
800°  the  product  is  in  the  form  of  a  powder,  but  at  a  some- 
what higher  temperature  this  melts  and  forms  a  hard  solid 
on  cooling.  This  oxide  is  used  in  making  several  lead 
salts  and  largely  in  the  manufacture  of  glass  of  certain 
kinds.  Red  lead,  Pb2O3  or  Pb3O4,  is  formed  by  heating 
litharge  in  the  air.  Oxygen  is  absorbed  to  form  what  may 
be  considered  as  a  mixture  of  PbO  and  PbO2.  It  is  used  as 
a  pigment  to  some  extent,  also  in  varnishes  and  in  the  glass 
industry  instead  of  litharge.  It  is  also  used  in  cements. 
Lead  peroxide,  PbO2,  is  made  by  action  of  nitric  acid  on  red 
lead,  by  which  means  the  PbO  is  dissolved  out.  It  may 
also  be  made  by  action  of  bleaching  powder  on  lead  chlo- 
ride and  is  valuable  as  an  oxidizing  agent  in  technology 
and  in  the  laboratory.  Several  hydroxides  are  known,  but 
they  are  not  specially  important. 

Lead  Sulphide  occurs  as  galena  and  may  be  obtained 
by  precipitation  of  a  lead  salt  with  hydrogen  sulphide. 
This  is  a  black  precipitate,  but  if  the  solution  contains 
much  free  hydrochloric  acid  a  reddish  chlorosulphide  is 
obtained. 


GENERAL  CHEMISTRY.  351 

Lead  Carbonates.  Several  are  known,  but  the  basic 
carbonate,  a  mixture  of  PbCO3  and  Pb(OH)2,  is  the  im- 
portant one  and  is  known  as  white  lead.  This  is  made  by 
a  number  of  processes,  but  the  so  called  Dutch  process 
yields  the  best  product.  In  this,  sheets  of  corrugated  lead 
are  placed  in  stoneware  jars  with  a  little  vinegar  and  these 
jars  are  imbedded  in  rows  in  spent  tanbark,  but  covered  so 
as  to  exclude  contamination.  The  spent  bark  evolves  CO2 
and  this,  acting  on  the  basic  acetate  of  lead  formed  in  the 
jars,  converts  it  into  basic  carbonate.  The  heat  generated 
by  the  putrefaction  of  the  bark  is  an  important  factor  in 
the  formation  of  the  basic  acetate.  The  whole  operation 
is  a  slow  one,  requiring  several  weeks  for  completion.  At 
the  end  of  this  time  the  jars  are  carefully  uncovered  and 
emptied  into  sifting  machines,  where  the  pure  white  lead  is 
separated  from  any  of  the  metal  not  corroded.  The  prod- 
uct is  ground  with  water  to  a  uniform  cream,  dried  and 
sold  in  this  form,  or  ground  in  oil.  Jt  is  the  most  valuable 
white  pigment  we  have,  but  is  often  adulterated  with 
cheaper  products. 

Lead  Chloride,  Bromide  and  Iodide  are  compounds 
which  are  soluble  in  boiling  water,  but  only  slightly  soluble 
in  cold  water.  Their  hot,  concentrated  solutions,  there- 
fore, deposit  crystalline  precipitates  on  cooling.  The  for- 
mulas are  PbCl2,  PbBr2  and  PbI2. 

Lead  Sulphate,  PbSO4,  is  a  white  substance,  insolu- 
ble in  water,  and  only  slightly  soluble  in  acids.  It  is  fre- 
quently used  as  a  pigment,  but  this  has  but  little  value. 

Lead  Nitrate  is  made  by  dissolving  litharge  in  nitric 
acid.  It  is  a  soluble  salt  and  is  used  in  making  other  lead 
compounds  by  precipitation,  and  also  as  a  mordant  by  the 
dyer. 

Lead  Chromate,  PbCrO4,  is  known  as  chrome  yel- 
low. It  is  made  by  precipitating  the  nitrate  or  acetate  by 
a  solution  of  potassium  dichromate  and  is  used  as  a  pig- 
ment. Chrome  orange  is  basic  lead  chromate. 

Lead  Acetate,  Pb(C2H3O2)2,  is  the  most  valuable  solu- 


352  GENERAL  CHEMISTRY. 

ble  compound  of  lead.  It  crystallizes  with  3H2O  and  so 
comes  into  commerce.  A  solution  of  this  salt  in  water 
has  the  power  of  dissolving  large  amounts  of  litharge, 
forming  solutions  of  basic  lead  acetate,  or  subacetate.  All 
of  them  have  important  applications. 

Recognition.  Lead  compounds  are  recognized  by  the 
white  precipitate  of  sulphate  formed  by  the  action  of  solu- 
ble sulphates,  by  the  black  sulphide  formed  in  their  solu- 
tions on  addition  of  hydrogen  sulphide  or  ammonium  sul- 
phide or  by  the  precipitation  of  white  lead  chloride  on 
addition  of  a  soluble  chloride.  This  last  precipitate  dis- 
solves by  eating,  but  comes  down  in  crystalline  needles 
on  cooling. 

TITANIUM. 

Although  not  abundant,  this  metal  is  widely  distrib- 
uted, occurring  as  dioxide,  TiO2,  in  rutile  and  two  other 
minerals,  and  also  with  oxides  of  iron  and  other  ores.  It 
is  therefore  a  constituent,  in  small  amount,  of  some  cast 
irons.  The  metal  maybe  obtained  in  the  free  state,  but  is 
not  technically  useful.  In  some  of  its  compounds  it  acts 
the  part  of  a  true  metal,  forming  well  defined  halogen  and 
other  salts,  while  in  other  compounds  it  acts  as  titanic  acid, 
yielding  titanates  resembling  the  stannates.  The  metal 
combines  directly  with  nitrogen  at  a  high  temperature. 

ZIRCONIUM. 

This  is  a  comparatively  rare  metal,  found  in  the  min- 
eral zircon,  ZrSiO4,  and  in  other  even  rarer  minerals.  The 
dioxide,  ZrO2,  becomes  incandescent  when  heated,  and  has 
been  used  in  making  lamps;  but  these  have  been  displaced 
largely  by  the  lamps  in  which  oxide  of  thorium  is 
employed. 

CERIUM. 

This  metal  occurs  in  a  few  rare  minerals,  of  which 
monazite,  a  complex  phosphate  of  cerium,  thorium  and 
lanthanum,  ccrite,  a  silicate  of  cerium,  lanthanum  and 


GENERAL  CHEMISTRY.  353 

other  metals,  and  orthite,  a  silicate  of  aluminum,  cerium, 
iron  and  calcium,  are  the  most  important.  The  oxides  of 
cerium  and  several  associated  metals  are  commonly  spoken 
of  as  the  cerite  earths.  Interest  attaches  to  these  bodies 
mainly  because  of  their  use  in  the  construction  of  lamps 
for  illumination  by  gas  by  incandescence.  One  of  the  salts 
of  cerium,  the  oxalate,  is  used  in  medicine. 

THORIUM. 

This  rare  metal  is  found  in  monazite,  mentioned  above, 
and  also  in  thorite  and  orangeite,  which  are  complex  silicates. 
The  metal  itself  has  no  technical  value,  but  its  dioxide, 
ThO2,  has  become  of  immense  importance  because  of  its 
power  of  emitting  a  brilliant  white  light  at  a  relatively  low 
temperature.  The  body  which  becomes  incandescent  is  a 
gauze  like  skeleton,  cone  or  mantle  of  thorium  dioxide 
mixed  with  a  small  amount  of  cerium  oxide  with  traces  often 
of  the  oxides  of  yttrium,  neodymium  and  zirconium.  To 
obtain  this  skeleton,  cotton  gauze  woven  or  cut  to  the  desired 
size  and  shape  is  immersed  in  a  solution  of  the  prepared 
thorium  nitrate.  It  is  then  dried  and  carefully  ignited,  by 
which  means  water,  oxides  of  nitrogen  and  organic  matter 
are  expelled,  leaving  the  metallic  oxides  in  the  shape  of  the 
original  cotton  gauze  skeleton.  When  this  is  properly 
supported  in  the  colorless  gas  flame,  that  of  a  Bunsen 
burner,  for  instance,  it  becomes  brightly  incandescent.  It 
has  been  shown  that  thorium  oxide  alone  does  not  yield 
the  brightest  light,  but  that  this  is  secured  by  the  presence 
of  the  cerium  oxide. 


CHAPTER    XIX. 


THE  NITROGEN  GROUP  OF  ELEMENTS  :  NITROGEN, 
PHOSPHORUS,  VANADIUM,  ARSENIC,  COLUM- 
BIUM,    ANTIMONY,    TANTALUM    AND 
BISMUTH. 

GENERAL  CHARACTERISTICS. 

T^HE  important  properties  of  some  of  the  elements  of 
1  this,  the  fifth  family  or  group  in  the  Periodic  System, 
have  been  discussed  already.  It  remains  to  point 
out  family  resemblances  here  and  to  describe  briefly 
the  compounds  of  the  remaining  important  elements.  In 
their  chemical  behavior  the  first  elements  in  the  family 
are  decidedly  acid,  while  as  before  observed  in  other 
groups,  with  increase  in  atomic  weight  the  elements 
become  decidedly  metallic.  The  physical  constants,  so 
far  as  known,  are  shown  in  the  table  below  : 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

Nitrogen  

14.04 

—200° 

Phosphorus 

31  02 

1  83 

16  9 

44  2° 

Vanadium 

51  38 

5  50 

9  3 

very   high 

Arsenic 

75  01 

5  73 

13  1 

500° 

Columbium   

93  73 

7  20 

13  0 

very  high 

Antimony   .... 

120  43 

6  71 

17  9 

440° 

Tantalum       

182  84 

10  60 

17  2 

very  high 

Bismuth  

208.11 

9  82 

21  2 

268° 

The  elements  of  the  group  are  in  most  of  their  com- 
pounds trivalent  or  pentavalent,  as  the  following  table 
illustrates  : 


GENERAL  CHEMISTRY.  355 

NH,        NC13  N203  N205  HN03 

PH,         PCI,  P203  PC15  P205  HP03 

VC13  V203  VA  HV03 

AsH3       AsCl3  As2O3  As2O5  HAsO3 

CbCl3  CbCl5         Cb,O5  HCbO3 

SbH3       SbCl3  Sb2O3  SbCl5         Sb2O5  HSbO3 

TaCl6         Ta2O5  HTaO3 

BiCl3  Bi2O3  Bi2O5  HBiO3 

The  parallelism  between  nitrogen  and  vanadium  is 
further  shown  by  the  series  of  oxides  formed,  these  five 
being  known:  V2O,  V2O2,  V2O3,  V2O4  and  V2O5. 

In  the  last  column  of  the  table  above  the  formulas  of 
the  meta  acids  corresponding  to  the  pentoxides  are  given. 
With  the  first  in  the  series,  HNO3,  the  acid  properties  are 
very  strong,  while  with  the  last,  HBiO3,  they  are  quite 
absent.  The  others  show  acid  behavior  decreasing  in 
character  as  we  pass  down  in  the  series.  Salts  of  all  these 
acids  are  known  until  we  reach  the  last. 


VANADIUM. 

This  rare  element  is  found  in  the  form  of  lead  and 
bismuth  salts  of  vanadic  acid  in  several  minerals.  Vana- 
dium changes  its  valence  or  capacity  for  holding  different 
amounts  of  oxygen  very  readily  and  on  account  of  this  the 
dioxide  has  found  a  very  important  application  in  the 
manufacture  of  anilin  black,  where  it  acts  as  a  carrier  of 
oxygen  or  strong  oxidizing  agent.  The  dioxide  becomes 
pentoxide,  which  is  in  turn  readily  reduced  by  giving  up 
its  oxygen  to  the  anilin  compound  undergoing  oxidation. 
One  part  of  the  dioxide  in  presence  of  potassium  chlorate 
is  able  to  convert  1,000  parts  of  anilin  hydrochloride  into 
anilin  black. 

COLUMBIUM  AND  TANTALUM. 

These  are  two  very  rare  metals  which  occur  usually 
together  in  isomorphous  minerals  known  as  columbite  and 
tantalite.  The  first  is  often  called  niobium. 


356  GENERAL  CHEMISTRY. 

ANTIMONY. 

Occurrence.  This  important  metal  is  found  in  nature 
mainly  as  the  sulphide,  or  stibnite,  Sb2S3.  Native  antimony 
occurs  in  small  quantities,  and  in  many  silver,  lead  and 
copper  ores  the  metal  is  found  in  combination,  usually  as 
sulphide. 

History.  The  gray  ore  or  stibnite  has  been  used  since 
a  remote  period  for  the  painting  of  the  eyebrows,  as 
described  by  Hebrew  and  Greek  writers.  The  separation 
of  the  metal  was  first  described  in  the  fifteenth  century,  and 
because  of  certain  virtues  it  was  supposed  to  possess  was 
very  thoroughly  studied.  Antimony  played  an  important 
part  in  the  search  for  the  philosopher's  stone  and  in  the 
attempts  at  transmutation  of  metals. 

Metallurgy.  Two  processes  are  in  use  for  the  produc- 
tion of  the  metal.  In  one  the  ore  is  roasted  to  form  the 
oxide,  Sb2O3,  which  is  reduced  then  by  heating  with  carbon: 


In  the  second  process  the  sulphide  is  melted  in  plum- 
bago crucibles  with  scrap  iron.  A  slag  of  ferrous  sulphide 
separates  from  the  reduced  heavier  metallic  antimony. 
Sometimes  the  ore  is  melted  in  crucibles  with  a  mixture  of 
ferric  oxide  and  charcoal,  which  accomplishes  the  same 
result.  Metallic  antimony  is  produced  in  Germany,  Aus- 
tria and  Hungary,  France,  Italy  and  the  United  States. 
The  production  of  this  country  in  1896  was  613  tons, 
mostly  from  foreign  ores.  The  world's  production  was 
about  3,500  tons. 

Properties.  Antimony  is  a  light  colored  crystalline 
and  very  brittle  metal  which  does  not  oxidize  at  ordinary 
temperatures,  but  which  combines  easily  with  chlorine  or 
bromine.  When  heated  it  burns  in  the  air.  When  heated 
with  nitric  acid  it  becomes  oxidized  to  pentoxide.  Cold, 
dilute  hydrochloric  and  sulphuric  acids  have  but  little  ac- 
tion on  it.  It  combines  readily  with  several  metals,  form- 
ing valuable  alloys. 


GENERAL  CHEMISTRY.  357 

Uses.  It  is  employed  in  making  a  few  compounds, 
but  principally  in  the  production  of  several  alloys.  Com- 
mon type  metal  is  an  alloy  of  about  4  parts  of  lead  to  1  of 
antimony.  Some  very  hard  type  metals  contain  50  parts 
of  lead,  25  parts  of  antimony  and  25  parts  of  tin.  Bri- 
tannia metal  contains  80  parts  or  more  of  tin,  10  to  15  of 
antimony  and  small  amounts  of  zinc  or  copper.  White 
metal,  used  for  bearings,  contains  80  to  90  parts  of  tin,  7 
to  15  of  antimony  and  2  to  10  of  copper.  Babbitt  metal 
contains  tin,  lead,  antimony  and  copper  in  different  pro- 
portions for  different  purposes. 

Antimony  Oxides.  Three  of  these  bodies  are  known, 
Sb2O3,  Sb2O4  and  Sb2O5.  The  first  is  formed  by  oxidiz- 
ing antimony  with  dilute  nitric  acid  or  by  decomposing 
the  chloride  by  a  solution  of  sodium  carbonate: 

2SbCl3+3Na2CO3  =  Sb2O3+6NaCl-f3CO2. 

It  is  a  white  powder  insoluble  in  water  but  soluble  in 
hydrochloric  acid  and  forming  SbCl3.  This  oxide  also  dis- 
solves in  a  solution  of  cream  of  tartar,  yielding  potassium- 
antimony  tartrate,  or  tartar  emetic,  2(KSbOC4H4O6).H2O. 

The  tetroxide,  Sb2O4,  is  formed  by  heating  either  the 
trioxide  or  pentoxide  in  the  air;  one  takes  up  oxygen,  the 
other  loses.  This  oxide  is  therefore  the  most  stable  of 
the  three.  The  pentoxide,  Sb2O5,  is  made  by  oxidizing  the 
metal  with  strong  nitric  acid.  It  is  a  light  yellow  powder, 
but  slightly  soluble  in  water,  but  which  behaves  as  an  acid 
anhydride  with  alkalies.  When  fused  with  the  latter  it 
yields  antimonates.  From  some  of  these  salts  antimonic 
acid,  HSbO3,  may  be  separated  as  a  white  powder,  but 
slightly  soluble  in  water.  The  resemblance  of  antimony 
compounds  to  those  of  phosphorus  is  shown  by  the  follow- 
ing formulas: 

HP03       HSbO3      NaSb03         NH4SbO3  KSbO3 

H3P04    H3Sb04 

H4P207  H4Sb207   H2Na2Sb2O7  (NH4)4Sb2O7    K4Sb2Ov 

The  salt  H2Na2Sb2O7,  or  acid  sodium  pyroantimonate, 
is  the  most  insoluble  salt  of  this  metal  known  and  is 


358  GENERAL  CHEMISTRY. 

obtained  in  analytical   chemistry  as  a  precipitate   in  the 
detection  of  sodium. 

Antimony  Sulphides.  The  trisulphide,  Sb2S3,  and 
the  pentasulphide,  Sb2S5,  are  known.  The  first  occurs  in 
nature  as  stibnite,  and  may  be  made  by  precipitation  of  a 
solution  of  antimony  trichloride  by  H2S.  So  obtained  it  is 
orange  yellow  and  amorphous.  This  sulphide  dissolves  in 
an  excess  of  hydrochloric  acid  and  is  soluble  also  in  alkali 
sulphides  forming  sulphantimonites: 

3(NH4)2S+Sb2S3  =  2(NH4)3SbS3. 

This  reaction  is  of  great  importance  in  analytical  chem- 
istry. 

The  pentasulphide,  Sb2S6,  is  a  dark  orange  precipitate 
formed  by  adding  hydrogen  sulphide  to  a  solution  of  the 
pentachloride.  It  dissolves  in  alkali  sulphides  to  form 
sulphantimonates : 

Sb2S5+6NaSH  =  2Na3SbS4+3H2S. 

The  salt,  Na3SbS4.9H2O,  is  known  as  Schlippe's  salt. 
The  pentasulphide  is  prepared  in  quantity  by  boiling  a 
mixture  of  powdered  stibnite,  sulphur  and  caustic  soda 
and  decomposing  the  Na3SbS4  formed  by  means  of  hydro- 
chloric acid.  It  is  employed  in  vulcanizing  rubber.  Several 
oxysulphides  of  antimony  are  known,  which  at  one  time 
were  largely  used  in  medicine. 

Antimony  Chlorides.  The  trichloride,  SbCl3,  is  formed 
as  a  soft  mass  by  passing  chlorine  over  antimony,  the 
metal  being  in  excess.  This  chloride  melts  easily  and  boils 
at  220°.  With  water  it  decomposes,  yielding  a  white  pre- 
cipitate, SbOCl,  known  as  the  powder  of  Algaroth. 

The  pentachloride,  SbCl5,  is  a  colorless  liquid  ob- 
tained by  passing  chlorine  in  excess  over  antimony.  When 
heated  it  decomposes  into  the  trichloride  and  free  chlorine. 
The  behavior  of  the  metal  with  chlorine  has  been  illustra- 
ted in  an  earlier  chapter. 

Antimony  and  Hydrogen.  The  hydride  SbH3,  a 
gaseous  body,  is  easily  formed  by  the  action  of  nascent 


GENERAL  CHEMISTRY,  359 

hydrogen  on  antimony  compounds.  The  method  em- 
ployed for  the  preparation  of  the  analogous  arsenic  com- 
pound, AsH3,  described  in  Chapter  IX,  may  be  used  here. 
The  antimony  compound  is  called  stibine.  It  yields  a  dark 
stain  on  porcelain  very  similar  to  that  from  arsenic,  but 
this  stain  is  not  soluble  in  hypochlorite  solutions,  while 
that  of  arsenic  is. 

Recognition.  Antimony  compound's  are  best  recog- 
nized through  the  formation  and  decomposition  of  SbH3, 
also  by  the  orange  yellow  precipitates  their  solutions  give 
with  hydrogen  sulphide,  and  the  solubility  of  these  pre- 
cipitates in  alkalies  or  alkali  sulphide  solutions. 


BISMUTH. 

Occurrence.  This  metal  is  found  in  the  native  condi- 
tion usually,  and  sometimes  as  the  oxide,  Bi2O3,  the  sul- 
phide, Bi2S3,  and  in  a  few  rare  minerals.  Mines  in  Saxony 
furnish  the  largest  part  of  the  metal. 

History.  Bismuth  has  been  known  since  the  thirteenth 
century,  but  was  often  confounded  with  antimony  and  tin. 
Its  true  nature  and  important  properties  were  pointed  out 
about  the  middle  of  the  last  century. 

Metallurgy.  The  ores  are  roasted  and  melted  in  pots 
with  iron  and  carbon  yielding  a  crude  bismuth  containing 
traces  of  other  metals.  To  make  the  pure  metal  the  crude 
ingots  are  heated  over  a  low  fire  on  inclined  iron  plates. 
The  pure  bismuth  melts  and  flows  out,  leaving  the  impuri- 
ties behind.  Bismuth  ores  often  occur  mixed  with  cobalt 
ores.  In  such  a  case  slag  from  a  previous  operation  is 
added  to  the  iron  and  carbon  so  that  cobalt  collects  in  the 
new  slag  and  is  worked  up  as  smalt. 

The  metal  is  produced  chiefly  by  two  works  in  Saxony, 
and  by  the  English  firm  of  Johnson,  Matthey  &  Co.,  from 
Australian  and  Bolivian  ores. 

Properties.  The  metal  is  hard  and  very  brittle,  with 
a  relatively  low  melting  point.  It  has  a  peculiar,  reddish 


360  GENERAL  CHEMISTRY. 

white  color  which  is  characteristic.  At  a  red  heat  bismuth 
decomposes  water;  it  combines  directly  with  the  halogens, 
but  is  acted  on  only  slowly  by  hydrochloric  or  sulphuric 
acid.  Nitric  acid  is  the  best  solvent.  It  alloys  with  the 
heavy  metals,  being  especially  valuable  in  the  production 
of  fusible  alloys. 

Uses.  It  is  mainly  employed  in  making  alloys.  Rose's 
metal  contains  2  parts  of  bismuth,  1  part  of  lead  and  1  part 
of  tin,  melting  at  94°.  Wood's  metal,  melting  at  61°,  con- 
tains 4  parts  of  bismuth,  2  parts  of  lead,  1  part  of  tin  and  1 
part  of  cadmium.  It  is  possible  to  make  mixtures  which 
melt  at  other  temperatures  between  65°  and  150°. 

These  alloys  are  employed  as  soft  solders  and  very 
largely  in  making  fusible  safety  plugs  for  steam  boilers, 
as  they  may  be  made  to  melt  at  some  definite  temperature 
corresponding  to  a  pressure  which  for  safety  should  not 
be  exceeded.  Water  pipes  in  automatic  sprinkling  sys- 
tems, for  protection  against  fire,  are  closed  with  plugs  with 
low  melting  points.  If  the  pipes  become  warm  the  plugs 
melt  and  allow  the  water  to  escape  in  such  a  manner 
as  to  extinguish  the  fire.  Fusible  alloys  are  often 
employed  in  making  type  for  temporary  use  in  newspaper 
printing;  this  is  possible  as  they  expand  slightly  on  cool- 
ing, thus  giving  a  sharp  impression. 

Bismuth  is  used  in  making  several  compounds,  which 
are  employed  as  cosmetics  and  in  medicine. 

Bismuth  Oxides.  Four  are  known,  Bi2O2,  Bi2O3,  Bi2O4 
and  Bi2O5.  The  trioxide  is  the  most  important;  it  is  a  yel- 
lowish powder  and  is  sometimes  used  in  making  certain 
kinds  of  glass. 

Halogen  Compounds.  BiCl3,  BiBr3and  BiI3  are  well 
known.  The  first  in  pure  condition  is  a  syrupy  liquid 
which  decomposes  with  water  yielding  the  oxychloride, 
BiOCl.  The  bromine  and  iodine  compounds  decompose 
also,  but  not  so  readily. 

Bismuth  Nitrates.  The  metal  or  the  oxide  may  be 
dissolved  in  nitric  acid  yielding  the  trinitrate,  Bi(NO3)3, 


GENERAL  CHEMISTRY.  361 

which  may  be  crystallized  from  the  acid  solution.  When 
mixed  with  water  this  nitrate  decomposes,  forming  a  basic 
nitrate: 

=  Bi(OH)2NO3-f2HNO3. 


This  basic  nitrate,  or  subnitrate,  is  an  important  medic- 
inal substance. 

Other  Compounds.  Sulphides,  a  basic  carbonate  and 
several  chromates  are  among  the  other  well  known  bis- 
muth compounds.  They  are  not  important,  however. 

Recognition.  Bismuth  compounds  are  not  stable 
except  in  presence  of  an  excess  of  acid.  Therefore  pre- 
cipitation follows  when  the  acid  solutions  are  mixed  with 
a  large  excess  of  water,  and  this  constitutes  a  very  delicate 
test  for  them.  A  black  sulphide  is  formed  by  precipitat- 
ing a  hydrochloric  acid  solution  with  hydrogen  sulphide. 


CHAPTER  XX. 


THE    CHROMIUM    GROUP:     CHROMIUM,    MOLYBDE- 
NUM, TUNGSTEN  AND  URANIUM.— RELA- 
TIONS TO  THE  OXYGEN  GROUP. 

GENERAL  CHARACTERISTICS. 

THESE  elements  form  the  primary  and  secondary  groups 
of  the  sixth  family  in  the  Periodic  System.  In  some 
respects  the  two  groups  show  little  in  common,  while 
from  other  standpoints  many  important  likenesses  may  be 
discovered,  some  of  which  will  be  shown  below.  The  ele- 
ments will  first  be  arranged  in  the  order  of  their  atomic 
weights  in  tabular  form. 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

BOILING 
POINT. 

16  00 

—180° 

Sulphur               

32.07 

2.05 

15.7 

114° 

440° 

Chromium              

52.14 

6.7 

7.7 

high 

Selenium           

79.02. 

4.60 

17.1 

217° 

665° 

Molybdenum      

95.98 

8.60 

11.1 

high 

Tellurium   

127.49 

6.25 

20.4 

500° 

abovel.OOO0 

Tungsten 

184  83 

19  1 

9.7 

high 

Uranium  

239.59 

18.7 

12.8 

high 

Sulphur,  selenium  and  tellurium  are  distinctly  non- 
metallic  elements,  and  this  is  shown  in  most  of  their  com- 
pounds. In  physical  behavior  chromium,  molybdenum, 
tungsten  and  uranium  are  characteristically  metallic  and 
in  some  of  their  compounds  this  is  plainly  shown.  In 
most  of  them,  however,  the  acid  character  is  very  pro- 
nounced. As  metals,  chromium  and  uranium  form  a  num- 


GENERAL  CHEMISTRY.  3G3 

her  of  important  salts,  while  tungsten  and  molybdenum 
are  important  in  the  compounds  they  form  with  metals. 
The  elements  of  this  family  form  acid  oxides  of  the  type 
SO3,  to  which  acids  and  well-known  salts  correspond. 

SO3  H2SO4  CrO3  H2CrO4 


(Se03)  H2Se04 

Te03  H2Te04 


MoO3          H2MoO4 
WO3  H2WO4 

U03  (H2U04) 


CHROMIUM. 

Occurrence.  The  metal  is  found  principally  as 
chrome  ironstone,  Cr2O3.FeO.  It  also  occurs  in  lead 
chromate  or  crocoisite,  PbCrO4.  In  traces  it  occurs  in 
several  minerals,  to  which  it  imparts  color,  the  emerald, 
for  illustration. 

History  and  Preparation.  It  was  discovered  in  1797 
in  the  investigation  of  lead  chromate,  and  in  the  same  year 
its  presence  was  recognized  in  the  mineral  called  chrome 
ironstone.  The  separation  of  the  metal  is  somewhat 
difficult.  It  has  been  made  by  the  reduction  of  the  oxide, 
Cr2O3,  at  a  very  high  temperature  by  aid  of  carbon,  and 
also  by  action  of  metallic  sodium  on  the  chloride,  CrCl3, 
at  a  high  temperature.  The  amount  of  chromium  pro- 
duced at  the  present  time  is  quite  small. 

Properties  and  Uses.  Chromium  is  a  bright,  light 
colored  metal,  extremely  hard  and  very  difficult  to  fuse. 
It  forms  a  hard  alloy  with  iron  known  as  chrome  steel, 
which  possesses  very  valuable  properties,  and  which  may 
find  important  applications  in  the  future  when  methods  of 
producing  it  are  better  understood. 

Chromium  Oxides.  Several  are  known,  of  which  the 
sesquioxide,  Cr2O3,  and  the  trioxide,  CrO3,  are  important. 
The  sesquioxide  is  known  as  chrome  green,  and  is  made  by 
igniting  the  hydroxide  described  below,  or  by  strongly 
heating  a  mixture  of  potassium  dichromate,  sodium  car- 
bonate and  ammonium  chloride.  On  extracting  the  mass 


364  GENERAL  CHEMISTRY. 

formed  with  water  the  green  oxide  is  left.     It  is  used  as  a 
valuable  pigment. 

The  formation  of  the  green  oxide  is  shown  by  this 
experiment  : 

Ex.  169.  Mix  10  Gm.  of  powdered  potassium  dichromate  with  an 
equal  weight  of  ammonium  chloride  and  a  gram  of  dry  sodium  carbo- 
nate. Rub  the  mixture  in  a  mortar  and  then  partly  fill  a  test-tube  with 
it.  Heat  the  test-tube  to  a  high  temperature  until  the  whole  mass  appears 
dark  green.  Then  dip  the  tube  in  cold  water  in  a  beaker.  This  breaks 
the  glass  and  allows  the  contents  to  mix  with  the  water.  The  green 
oxide  remains  suspended  while  the  other  matters  go  into  solution.  This 
reaction  amounts  to  a  reduction  of  the  dichromate. 

The  trioxide,  CrO3,  is  made  in  the  form  of  bright  red 
needles  by  the  action  of  strong  sulphuric  acid  on  potassium 
dichromate.  It  is  often  spoken  of  as  chromic  acid  or 
chromic  anhydride,  and  is  employed  because  of  its  strong 
oxidizing  properties.  It  is  used  in  laboratories  on  this 
account,  in  place  of  the  dichromate  which  acts  in  a  similar 
manner,  as  will  be  shown  below. 

Chromium  Hydroxides.  Of  these  the  best  known  is 
the  compound,  Cr(OH)3,  made  by  precipitating  a  chromic 
salt  with  ammonia  water.  It  is  a  bulky,  greenish  blue 
precipitate,  which  when  dried  and  ignited  yields  Cr2O3. 

Halogen  Compounds.  The  chloride,  CrCl3,  is  the 
most  important.  It  may  be  made  by  passing  a  current  of 
dry  chlorine  over  a  mixture  of  the  oxide,  Cr2O3,  and  car- 
bon, and  appears  in  the  form  of  pinkish  scales.  A  solution 
of  the  chloride  is  best  made  by  boiling  the  trioxide  with 
hydrochloric  acid.  A  dark  red,  fuming  liquid,  known  as 
chromium  oxychloride,  CrO2Cl2,  is  made  by  distilling  a 
mixture  of  potassium  dichromate,  sulphuric  acid  and  com- 
mon salt.  It  boils  at  118°  and  is  decomposed  on  contact 
with  water. 

Chrome  Alum,  or  chromium-potassium  sulphate, 
K2SO4.  Cr2(SO4)3. 24H2O,  is  a  very  common  substance  made 
from  potassium  dichromate.  At  the  present  time  it  is  pro- 
duced in  great  quantities  as  a  by-product  in  certain  indus- 
tries in  which  the  dichromate  is  employed  as  an  oxidizing 
agent.  An  illustration  will  be  given  of  this  in  the  following 


GENERAL  CHEMISTRY.  365 

experiment,  in   which   alcohol  is  converted  into   a  body 
known  as  aldehyde,  by  oxidation  : 

Ex.  170.  Dissolve  15  Gm.  of  potassium  dichromate  in  50  Cc.  of 
water  and  add  about  13  to  15  Cc.  of  strong  sulphuric  acid.  Cool  the 
mixture,  and  pour  into  it  10  Cc.  of  alcohol.  The  color  changes  from  red 
to  dark  violet  or  olive  green;  and  a  vapor,  with  a  sharp,  characteristic 
odor,  escapes.  This  is  aldehyde.  Allow  the  liquid  to  evaporate  spontane- 
ously, which  will  require  several  days,  and  observe  the  crystals  of 
chrome  alum  formed. 

The  color  of  the  solution  of  chrome  alum  depends  on 
its  temperature.  If  made  at  a  low  temperature  it  is  dark 
violet,  but  this  turns  to  green  by  application  of  heat,  from 
the  formation  of  complex  salts  in  which  the  chromium  and 
acid  are  differently  united. 

The  reaction  in  the  above  experiment  appears  complex, 
but  may  be  shown  to  be  this  : 

K2Cr207+4H2S04+3C2H60+24H20  = 

K2SO4.Cr2(SO4)3.24H2O+3C2H4O4-7H2O. 

Each  molecule  of  alcohol  has  lost  two  atoms  of  hydro- 
gen, and  this  hydrogen,  plus  that  from  the  sulphuric  acid, 
has  been  taken  by  the  seven  atoms  of  oxygen  in  the 
dichromate  to  form  seven  molecules  of  water. 

A  solution  of  chrome  alum  yields  a  precipitate  with 
ammonia  water,  this  precipitate  consisting  of  chromium 
hydroxide,  Cr(OH)3,  while  a  solution  of  the  dichromate 
yields  no  precipitate  with  alkalies,  but  changes  color  from 
formation  of  chromate,  as  the  following  equations  illus- 
trate : 


K2S04.Cr2(S04)3+6NH4OH  = 

2Cr(OH)3+K2S04+3(NH4)2S04. 

The  Chromates.  These  are  the  compounds  of  chro- 
mium trioxide,  and  include  several  extremely  important 
substances  of  which  sodium  and  potassium  dichromates 
are  the  best  known.  These  are  prepared  indirectly  from 
chrome  ironstone.  The  ore  is  ground  and  heated  to  a 
high  temperature  with  lime  or  chalk,  the  mass  being  well 
stirred  to  give  the  air  access. 

4Cr2O3FeO+7O2+8CaCO3=8CaCrO4+8CO2+2Fe2O3. 


366  GENERAL  CHEMISTRY. 

The  cool  mass  is  then  extracted  with  water  and  the 
calcium  chromate  solution  thus  obtained  is  converted  into 
the  sodium  or  potassium  compound  by  double  decomposi- 
tion with  an  alkali  sulphate.  The  alkali  chromates  are  in 
turn  treated  with  acid  to  form  dichromates,  which  being 
less  soluble  may  be  more  easily  purified  by  crystallization: 


Solutions  of  the  chromates  are  yellow,  while  those  of 
the  dichromates  are  reddish  yellow.  Potassium  dichro- 
mate  was  for  a  long  time  the  most  important  salt  of  the 
class,  but  at  the  present  time  the  corresponding  sodium 
compound  is  very  largely  used,  because  of  its  greater  sol- 
ubility. The  chromates  of  the  heavy  metals  are  insoluble 
precipitates.  Lead  chromate,  PbCrO4,  is  known  as  chrome 
yellow,  and  is  made  by  precipitating  a  solution  of  the  di- 
chromate  with  lead  acetate.  Chrome  orange  is  a  basic  lead 
chromate,  PbCrO4.Pb(OH)2,  made  by  precipitating  an  al- 
kaline chromate  by  a  soluble  lead  salt.  Both  lead  chro- 
mates are  used  in  dyeing.  Silver  chromate,  Ag2CrO4,  is  a 
deep  red  precipitate  formed  on  addition  of  a  chromate  to 
silver  nitrate  solution.  Barium  chromate  appears  as  a  yel- 
lowish precipitate  on  mixing  chromate  solutions  with  those 
of  soluble  barium  compounds. 

Recognition.  Salts  in  which  the  chromium  acts  as  the 
metal  are  usually  pink  or  violet  and  give  down  the  chro- 
mium in  the  form  of  hydroxide  on  addition  of  ammonia. 
These  compounds  when  fused  with  alkaline  oxidizing 
agents  are  converted  into  yellow  salts  of  chromic  acid. 
These  salts  yield  the  precipitates  with  lead  and  barium 
compounds  mentioned  above. 

MOLYBDENUM. 

This  is  a  comparatively  rare  metal  found  mainly  in  the 
disulphide,  MoS2,  and  in  a  lead  combination,  PbMoO4. 
The  free  metal  has  no  technical  uses.  An  important  com- 
pound, the  trioxide,  MoO3,  is  made  from  the  sulphide  by 
oxidizing  it  in  the  air.  This  trioxide  is  technically  called 


GENERAL  CHEMISTRY.  367 

molybdic  acid.  It  forms  salts  called  molybdates,  of  which 
ammonium  molybdate  (NH4)2MoO4,  is  the  most  important. 
This  molybdate  is  a  useful  laboratory  reagent,  being  em- 
ployed in  the  detection  of  phosphorus  in  phosphates. 
Molybdic  acid  forms  very  complex  salts,  and  the  commer- 
cial ammonium  salt  has  a  formula  which  is  usually  given 
as  (NH4)6Mo7O21.24H:2p. 

A  solution  of  the  trioxide  in  sulphuric  acid  is  a  common 
reagent  employed  in  the  detection  of  alkaloids. 


TUNGSTEN. 

This  metal  is  found  commonly  in  wolframite,  FeWO4, 
and  scheelite,  CaWO4.  Also  in  several  rarer  minerals.  The 
metal  may  be  obtained  by  reduction  of  the  oxides  by 
hydrogen,  but  has  few  applications  in  the  arts.  It  is  some- 
times combined  with  iron  to  form  a  kind  of  steel.  One  of 
the  most  important  compounds  is  the  trioxide,  WO3,  which 
is  often  called  tungstic  acid.  It  may  be  obtained  by  decom- 
posing a  tungstate  with  an  acid  as  a  bulky  precipitate 
which  yields  a  yellow  powder  on  drying.  Sodium 
tungstate  may  be  made  from  the  native  wolframite  by  fus- 
ing the  latter  with  a  mixture  of  sodium  carbonate  and 
nitrate.  The  fused  mass  is  leached  with  water  and  the 
solution  so  obtained  is  evaporated  to  crystallization.  The 
crystals  have  not  the  simple  formula,  Na2WO4,  but  are 
much  more  complex,  resembling  in  this  respect  those  of 
molybdic  acid.  The  common  commercial  sodium  tungstate 
is  Na10W12O41  with  2lH2O  to  28H2O,  the  amount  of  water 
depending  on  the  temperature  of  crystallization.  A  solu- 
tion of  this  salt  is  used  as  a  mordant  in  dyeing  and  also  in 
impregnating  muslin  to  render  it  uninflammable.  It  is 
employed  in  the  preparation  of  several  reagents  for  labo- 
ratory purposes. 

Wolfram  Bronzes.  A  very  peculiar  class  of  com- 
pounds is  obtained  by  the  partial  reduction  of  tungstates 
by  heating  in  hydrogen,  illuminating  gas,  or  with  certain 
metals.  From  their  colors  and  physical  properties  these 
bodies  are  called  wolfram  or  tungsten  bronzes.  Potassium- 


368  GENERAL  CHEMISTRY. 

tungsten  bronze  is  a  purple  blue  powder,  having  the  for- 
mula K2W4O12-  A  yellow  sodium  tungsten  bronze*  has 
the  formula  Na5W6O18;  a  blue  bronze  the  formula 
Na2W5Oi5,  and  a  purple  red  bronze  the  formula  Na.2W3O9. 
These  are  combinations  of  varying  amounts  of  WO3  with 
alkali,  and  are  practically  used  as  bronze  paints  for  many 
purposes.  Many  complex  compounds  of  tungstic  acid 
with  boric  and  phosphoric  acids  are  known,  but  they  have 
no  technical  importance. 

URANIUM. 

This  rare  metal  is  found  in  a  few  ores,  the  most  impor- 
tant being  a  complex  oxide  known  as  pitch-blende.  The 
free  element  maybe  obtained  as  a  steel  white  substance  by 
reducing  one  of  the  chlorides  with  sodium.  By  fusing 
pitch  blende  with  a  mixture  of  sodium  carbonate  and  ni- 
trate a  soluble  sodium  uranate,  Na2U2O7,  is  obtained, 
which  is  known  as  uranium  yellow  and  which  is  employed 
in  coloring  glass.  By  double  decomposition  with  ammo- 
nium chloride  this  salt  yields  the  ammonium  compound, 
which  is  decomposed  by  heat  with  formation  of  the  oxides 
UO2  and  UO3.  From  this  mixture  several  other  uranium 
compounds  may  be  made.  The  most  important  of  these  is 
the  nitrate,  UO8  (NO3)2.6H2O,  which  is  obtained  in 
beautiful  yellow  fluorescent  crystals.  A  solution  of  this 
salt  or  the  corresponding  acetate  is  often  employed  in  the 
laboratory  for  the  determination  of  phosphates. 

In  the  formation  of  salts  uranium  seems  to  act  with  a 
valence  of  four  and  six,  giving  rise  to  the  uranous  and 
uranic  compounds.  As  illustrations  of  the  uranous  com- 
pounds we  have  the  dioxide,  UO2,  the  tetrachloride, 
UC14,  and  the  sulphate,  U(SO4)2.  Among  the  uranic 
compounds  we  have  the  trioxide,  UO3  or  (UO2)O,  the 
group  (UO2)  being  called  uranyl;  the  chloride  or  uranyl 
chloride,  UO2C12,  uranyl  sulphate,  UO2(SO4).3H2O,  and 
nitrate  referred  to  above,  UO2(NO3)2.6H2O.  These 
uranic  compounds  may  be  regarded  as  containing  the 
group  (UO2). 


CHAPTER  XXI. 


MANGANESE  AND  ITS  RELATIONS  TO   THE    HALO= 
GEN  GROUP. 

GENERAL  CHARACTERISTICS. 

IN  THE  seventh  family  of  the  Periodic  System  there  are 
two  groups  which  appear  to  have  but  little  in  common. 
In  the  primary  group  we  find  manganese  alone,  and  in  the 
secondary  group  the  halogen  elements  already  described. 
In  many  respects  manganese  resembles  iron  and  in  some 
of  the  salts  a  resemblance  is  found.  But  there  are  other 
considerations  which  justify  the  grouping  of  this  metallic 
element  with  such  pronounced  nonmetals  as  chlorine  and 
bromine.  Like  chromium  and  other  elements  described  in 
the  last  few  chapters,  manganese  forms  two  classes  of  com- 
pounds. In  some  it  acts  as  a  metal,  while  in  others  it  be- 
haves as  the  acid  element.  It  is  among  these  compounds 
that  we  must  look  for  the  analogies  with  the  chlorine  group. 
Manganese  forms  at  least  five  oxides,  the  highest  one  of 
which  is  Mn2O7.  This  with  water  yields  an  acid  with  the 
strongest  oxidizing  properties,  and  having  the  composition 
HMnO4.  The  halogen  oxides  of  the  type  X2O7  are  not 
known  but  the  acids  corresponding  are,  and  also  salts. 
We  have 

HC104  KC104          Ba(C104)2 

HBrO4  KBrO4          Ba(BrO4)a 

HI04  KI04  Ba(I04)2 

Mn2O7  HMnO4          KMnO4        Ba(MnO4)2. 

All  of  these  acids  may  be  looked  upon  as  related  to  the 
oxide  in  this  way: 

X207+H20=:2HX04. 


370  GENERAL  CHEMISTRY. 

The  most  important  compound  of  manganese  is  the 
salt  KMnO4.  This  is  isomorphous  with  the  perchlorate, 
with  which  it  will  crystallize  in  all  proportions.  In  connec- 
tion with  somewhat  analogous  chemical  behavior  this  rela- 
tion is  important  in  suggesting  the  place  of  manganese 
among  the  elements.  On  the  other  hand,  as  will  appear 
below,  manganese  forms  salts  which  in  structure  resemble 
the  chromates.  Thus  we  have  Na^MnO^  as  we  have 
Na2CrO4,  but  the  manganate  is  far  less  stable  than  is  the 
chromate,  and  the  corresponding  oxide,  MnO3,  is  known 
only  as  a  compound  of  little  stability. 

Occurrence.  Manganese  is  found  principally  in  the 
dioxide  known  as  pyrolusite,  MnO2,  in  braunite, -Mn2O3,  in 
hausmannite,  Mn3O4,  and  in  sulphides  and  rarer  ores. 
Many  iron  ores  contain  traces  of  manganese. 

Preparation  and  Properties.  Pure  manganese  is  not 
readily  obtained,  but  may  be  produced  as  is  iron  by  blast  fur- 
nace reduction,  employing  a  higher  temperature  and  more 
carbon.  This  yields  a  cast  manganese  containing  carbon. 
The  pure  metal  is  made  on  a  small  scale  by  reduction  of  the 
chloride  with  sodium  or  magnesium,  and  by  several  other 
processes. 

The  atomic  weight  of  manganese  is  54.99,  the  specific 
gravity  about  7.5,  with  the  specific  volume,  therefore,  7.3. 
The  pure  metal  is  very  hard  and  may  be  melted  only  at  a 
high  temperature.  It  oxidizes  very  readily  in  the  air  and 
decomposes  water  at  a  relatively  low  temperature. 

Uses.  In  pure  form  the  metal  has  no  technical  appli- 
cations. Several  alloys  are  important,  one  being  used  in 
the  production  of  Bessemer  steel. 

Manganese  Oxides.  At  least  ten  of  these  combina- 
tions have  been  described,  but  not  over  five  are  thoroughly 
well  known.  Manganous  oxide,  MnO,  is  a  greenish  or  gray 
powder  obtained  by  the  reduction  of  the  other  oxides  in  a 
current  of  hydrogen.  It  dissolves  in  acids  forming  mangan- 
ous  salts.  Manganic  oxide,  Mn2O3,  is  found  in  nature  as 
braunite,  and  may  be  prepared  artificially  also.  In  compo- 


GENERAL  CHEMISTRY.  371 

sition  it  corresponds  to  chromic  oxide,  Cr2O3,  and  ferric 
oxide,  Fe2O3.  With  acids  it  yields  manganic  salts,  as 
MnCl3  and  Mn2(SO4)3.  Manganous-manganic  oxide,  Mn3O4, 
occurs  as  hausmannite,  and  may  be  made  by  heating  the 
dioxide  to  a  high  temperature. 

The  dioxide,  MnO2,  is  the  best  known  and  by  far  the 
most  important  of  the  manganese  oxides.  It  is  found  as 
pyrolusite  in  many  parts  of  the  world  and  has  several 
technical  uses.  It  is  employed  with  potassium  chlorate  in 
the  manufacture  of  oxygen;  it  is  used  for  the  decomposi- 
tion of  hydrochloric  acid  in  the  chlorine  industry,  and  is 
used  in  immense  quantities  in  the  production  of  spiegelei- 
sen  for  the  Bessemer  steel  industry.  It  is  also  employed 
in  considerable  quantities  for  the  production  of  colorless 
glass  from  common  materials  which  contain  iron.  When 
heated  to  a  high  temperature  it  gives  up  one-third  of  its 
oxygen  : 


The  trioxide,  MnO3,  has  been  described,  but  it  is  not 
stable  or  important.  The  same  may  be  said  of  several 
others.  The  heptoxide,  Mn2O7,  is  obtained  as  a  dark-green 
liquid  by  the  action  of  sulphuric  acid  on  potassium  per- 
manganate. This  liquid  is  somewhat  volatile  and  decom- 
poses easily  with  explosive  violence.  With  organic  sub- 
stances it  unites  immediately  with  explosive  oxidation. 
With  water  it  yields  permanganic  acid,  as  explained  above. 

Manganese  Hydroxides.  Several  are  known,  but  they 
have  little  importance.  Manganous  hydroxide,  Mn(OH)2, 
is  obtained  by  precipitating  manganous  salts  with  al- 
kali hydroxides.  It  is  a  light  colored,  flocculent  precipi- 
tate which  soon  oxidizes  in  the  air  to  higher  hydroxides, 
with  dark  color. 

Manganese  Chlorides.  The  most  important  one  of 
these  is  the  well-known  pink  salt,  MnCl2.4H2O,  which  may 
be  obtained  by  the  solution  of  manganese  or  its  sulphide 
in  hydrochloric  acid,  or  as  a  by-product  in  the  chlorine  in- 
dustry from  MnO2.  The  salt  is  very  soluble  in  water  and 


372  GENERAL  CHEMISTRY. 

may  be  crystallized  with  different  amounts  under  special 
conditions. 

In  the  production  of  chlorine  from  hydrochloric  acid  and 
manganese  dioxide  it  is  commonly  assumed  that  the  reac- 
tion takes  place  between  the  two  substances  in  this  way: 

MnO2+4HCl  =  MnCl2+Cl2+2H2O. 

It  is  likely,  however,  that  other  chlorides  are  formed 
in  the  operation  to  be  subsequently  decomposed  in  this 
manner: 

MnCl4=MnCl2+Cl2. 

The  trichloride,  MnCl3,  is  probably  formed  also  when 
the  acid  is  cold  and  strong,  but  both  this  and  the  tetra- 
chloride  break  up  readily,  liberating  chlorine  and  leaving 
the  dichloride. 

MnBr2  and  MnI2  are  salts  corresponding  to  manganous 
chloride,  but  they  are  not  important. 

Manganese  Sulphates.  The  most  important  is  the 
manganous  sulphate,  MnSO4,  which  may  be  obtained  crys- 
tallized with  3,  4,  5  or  7H.jO,  the  amount  of  water  held 
depending  on  the  concentration  and  temperature  of  the 
crystallizing  solution.  Manganous  sulphate  and  the 
alkali  sulphates  crystallize  together.  Manganic  sul- 
phate, Mn2(SO4)3,  is  known,  but  its  solution  is  not  very 
stable.  With  alkali  sulphate  solutions  it  yields  alums, 
K2SO4.Mn2(SO4)3.24H2O  and  (NH4)2SO4.Mn2(SO4)3.24H3O, 
being  well  known.  This  behavior  connects  manganese 
with  iron  on  the  one  hand  and  with  aluminum  and  chro- 
mium on  the  other. 

Many  other  compounds  of  metallic  manganese,  a  sul- 
phide, MnS,  a  carbonate,  MnCO3,  a  nitrate,  Mn(NO3)2,  etc., 
are  known,  but  they  are  not  practically  valuable.  The  most 
important  combinations  of  manganese  are  those  in  which  it 
exists  in  highly  oxidized  or  acid  condition,  as  in  the  man- 
ganates  and  permanganates. 

Potassium  Permanganate.  This  important  salt, 
which  consists  of  deep  purple  crystals,  may  be  looked  upon 


GENERAL  CHEMISTRY.  373 

as  the  potassium  compound  of  permanganic  acid,  in  turn 
derived  from  the  heptoxide,  Mn2O7: 


Practically  the  salt  is  produced  by  a  very  different  re- 
action. Manganese  dioxide  in  powder  is  heated  with  a 
mixture  of  potassium  hydroxide  and  an  oxidizer,  either  the 
nitrate  or  chlorate,  which  gives  rise  to  a  green  salt,  the 
manganate,  K2MnO4: 

3MnO2+6KOH  +  KClO3=3K2MnO4+3H2O-fKCl. 

By  lixiviating  with  water  and  passing  CO2  into  the 
solution  this  crude  manganate  is  decomposed  with  forma- 
tion of  permanganate  and  other  products: 

3K2MnO4+2CO2=2K3CO3+2KMnO4+MnO2. 

The  purple  solution  is  filtered  from  the  precipitated 
dioxide  through  marble  dust  and  is  concentrated  to  the 
point  where  crystallization  takes  place.  As  the  salt  is  not 
very  soluble  in  cold  water  it  is  easily  made  pure  in  this 
way.  By  using  chlorine  the  whole  of  the  manganate  may 
be  saved  as  permanganate,  and  this  is  done  on  the  techni- 
cal scale: 

2K2Mn04+Cl2  =  2KMn04+2KCl. 

A  crude  permanganate  of  sodium  or  potassium  is  made 
by  treating  the  corresponding  crude  manganate  solution 
with  a  little  dilute  sulphuric  acid.  The  solution  so  obtained 
is  used  directly  in  large  quantities  in  the  oxidation  of 
sewage  and  for  similar  purposes. 

Practically  all  the  uses  of  the  permanganate  depend  on 
its  power  of  readily  liberating  oxygen,  and  because  of  this 
behavior  it  is  largely  used  in  technical  and  analytical 
chemistry. 

In  acid  solution  the  permanganate  behaves  as  an  oxi- 
dizing agent,  as  illustrated  in  the  following  equations: 

2KMnO4+5C2O4H2-f3H2SO4  = 

2MnSO4-fK2SO4-|-10CO2+8H.2O. 

In  this  case  oxalic  acid  is  oxidized  to  water  and  carbon 


374  GENERAL  CHEMISTRY, 

dioxide,  and  as  the  reaction  is  a  perfectly  sharp  and  definite 
one  a  valuable  process  of  quantitative  volumetric  analysis 
is  based  on  it.  The  next  equation  illustrates  the  use  of 
the  permanganate  in  another  volumetric  process,  where 
ferrous  iron  is  oxidized  to  ferric  iron: 


5Fe2(SO4)3+K2SO4-f2MnSO4-f-8H2O. 

These  changes  depend  on  the  fact  that  two  molecules 
of  the  permanganate  with  acid  liberate  five  atoms  of  oxygen: 


In  several  very  important  practical  cases  the  perman- 
ganate is  used  as  an  oxidizer  in  neutral  or  alkaline  solu- 
tion, but  the  reaction  is  then  a  different  one,  and  may  be 
illustrated  as  follows: 


If  the  solution  is  sufficiently  alkaline  to  begin  with  the 
MnO2  does  not  precipitate,  but  dissolves  to  form  a  manga- 
nite.  In  this  case  two  molecules  of  permanganate  give  up 
three  atoms  of  oxygen  instead  of  five.  When  the  student 
takes  up  the  subject  of  volumetric  analysis  the  high  im- 
portance of  all  these  reactions  will  be  recognized.  The 
reactions  with  sulphurous  acid  and  hydrogen  dioxide  re- 
ferred to  in  former  chapters  take  place  according  to  these 
equations: 

2KMnO4+5H2SO3+3H2SO4  = 

K2SO4-f-2MnSO4-f5H2SO4+3H2O 
2KMnO4+5H2O2+3H2SO4=K2SO4+2MnSO4+8H2O+5O2. 

Recognition.  The  salts  and  oxides  of  manganese 
become  converted  into  manganates  when  fused  with  an 
alkali  and  a  little  potassium  chlorate.  The  green  color  is 
characteristic.  The  formation  of  a  flesh  colored  precip- 
itate of  sulphide  when  a  manganous  salt  is  treated  with 
ammonium  sulphide  solution  is  also  characteristic. 


CHAPTER  XXII. 


THE    IRON   GROUP:     IRON,    NICKEL    AND  COBALT. 


GENERAL  CHARACTERISTICS. 

THESE  three  elements  form  a  natural  period  in  the 
eighth  family  of  the  Periodic  System.  They  resemble 
each  other  as  metals  and  also  in  the  character  of  many  of 
their  compounds.  Thus,  they  form  similar  sulphates  and 
double  sulphates  with  members  of  the  alkali  group.  They 
form  oxides  of  the  types  MO  and  M2O3.  From  another 
standpoint,  however,  they  may  be  looked  upon  as  forming 
a  gradual  transition  between  manganese  in  the  seventh 
family  and  copper  in  the  first. 

Because  of  the  closeness  of  their  atomic  weights  the 
proper  classification  of  these  metals  is  a  matter  of  some 
uncertainty.  It  is  even  held  by  several  chemists  that  co- 
balt and  nickel  are  not  true  elements,  but  mixtures  of 
metals  yet  to  be  separated.  This  question  is  one  which 
cannot  be  discussed  in  a  book  for  beginners.  Accepting 
the  usual  views,  the  following  table  shows  the  important 
physical  constants  of  the  group: 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

Iron   

56.02 

7.8 

7.2 

1800° 

Nickel  
Cobalt.  . 

58.69 
58.93 

8.9 
8.6 

6.6 
6.8 

1500° 
1800° 

In   the   properties  of    the  metals    and  their   common 
salts  cobalt  stands  between  iron  and  nickel. 


376  GENERAL  CHEMISTRY. 

IRON. 

Occurrence.  Metallic  iron  is  occasionally  found  as  a 
natural  substance,  but  in  this  form  it  is  quite  unimportant. 
It  is  relatively  abundant  and  widely  distributed  in  a  num- 
ber of  minerals,  of  which  the  oxides,  sulphide  and  carbon- 
ate are  technically  valuable.  The  natural  oxides  which 
are  important  are  hematite,  Fe2O3,  which  occurs  in  varie- 
ties known  as  specular  iron  ore,  micaceous  ore,  red  ochre 
and  others;  limonite,  Fe2O3-j-water,  of  which  brown 
hematite,  bog  iron  ore  and  yellow  ochre  are  varieties; 
magnetite,  Fe3O4,  or  the  magnetic  oxide  of  iron.  Franklin- 
ite  is  a  ferric  oxide  in  which  some  of  the  iron  is  replaced 
by  zinc.  Ferrous  carbonate,  FeCO3,  occurs  as  spathic 
iron  ore;  an  impure  variety  is  known  as  clay  ironstone. 
The  native  sulphide  of  iron,  FeS2,  is  known  as  iron  pyrite 
and  is  valuable  for  its  sulphur  rather  than  for  its  iron. 

History.  The  metal  has  been  used  from  the  earliest 
historical  times.  Ores  were  reduced  by  the  ancient 
Hebrews,  Egyptians,  Greeks  and  Romans  by  methods 
which  were  doubtless  very  crude.  In  China  and  India  the 
metal  was  apparently  known  at  a  remote  period.  After 
their  conquest  of  Britain  the  Romans  began  working  de- 
posits of  iron  in  localities  which  still  yield  the  ore  in  quan- 
tity. The  history  of  improvements  in  the  smelting  of  ores 
on  the  large  scale  is  obscure,  but  the  greatest  advance  was 
made  in  the  introduction  of  the  blast  furnace,  which  was 
probably  as  late  as  the  sixteenth  century.  With  the  recog- 
nition of  the  chemical  differences  between  cast  iron, 
wrought  iron  and  steel,  the  development  of  a  rational 
metallurgy  of  iron  became  possible;  and  since  the  middle 
of  last  century  great  advances  have  been  made  in  all  civi- 
lized nations  in  the  smelting  and  working  of  the  metal. 

Metallurgy.  The  oxides,  or  carbonate  yielding  oxide, 
are  the  only  ores  of  importance  for  the  production  of 
metallic  iron.  The  reduction  of  the  oxide  to  the  state  of 
metal  depends  on  the  behavior  with  carbon  monoxide  at 
a  high  temperature;  like  carbon  itself  this  body  is  a  strong 


GENERAL  CHEMISTRY,  377 

reducing  agent  under  proper  conditions,  and   with  ferric 
oxide  brings  about  this  change  : 


Blast  Furnace. 

The  smelting  of  iron  ores  is  carried  out  in  a  blast 
furnace,  the  general  arrangement  of  which  is  shown  in  the 
accompanying  figure. 

Such  furnaces  are  from  60  to  100  feet  high  and  from  15 
to  20  feet  wide  in  the  widest  part.  At  the  bottom  of  the 
furnace  there  are  openings,  called  tuyeres,  through  which 
a  blast  of  heated  air  may  be  blown  in  under  considerable 
pressure.  After  being  thoroughly  heated  by  combustion 
of  fuel  the  furnace  is  charged  with  alternate  layers  of 
coke  (charcoal  or  anthracite  is  sometimes  used),  ore  and 
limestone  to  the  top.  The  oxygen  of  the  blast  forms 
CO2  at  the  base  of  the  furnace,  but  this  becomes  reduced 
to  CO  by  the  excess  of  carbon  a  short  distance  higher. 
This  CO  in  turn  acts  on  the  oxide  of  iron,  reducing  it 
as  explained  above.  The  reduced  iron  gradually  settles 
down  in  the  furnace  and  finally  melts,  collecting  at  last  in 
the  liquid  condition  in  a  kind  of  hearth  at  the  bottom. 
Meanwhile  more  ore,  fuel  and  limestone  are  added  at  the 
top,  the  process  being  kept  up  continuously.  From  time  to 
time,  usually  every  eight  or  twelve  hours,  the  hearth  below 
is  tapped  by  withdrawing  a  plug  and  the  molten  iron  col- 
lected is  allowed  to  run  out  into  trenches  made  in  sand  of 
the  casting  floor  in  front  of  the  furnace.  The  iron  soon 
solidifies  in  the  form  of  rough  bars,  which  are  technically 
called  "  pigs."  This  pig  iron  is  essentially  what  is  known 
as  cast  iron,  the  nature  of  which  will  be  soon  explained. 

As  all  ores  contain  certain  earthy  matters  it  is  neces- 
sary to  add  something  to  the  charge  in  the  furnaces 
to  separate  these  in  the  form  of  a  slag.  Limestone  is 
usually  added  because  the  impurities  are  commonly  sili- 
cates. The  limestone  is  called  a  flux,  as  it  forms  with  the 
silicate  an  impure  calcium  silicate  which  separates  as  a 
slag  or  glass  and  collects  finally  as  a  liquid  layer  just 
above  the  molten  iron.  Through  an  opening  above  that  for 


378 


GENERAL  CHEMISTRY. 


FIG.  33. 


GENERAL  CHEMISTRY.  379 

the  iron,  this  slag  is  run  off  from  time  to  time.  In  settling 
down  toward  the  hearth  of  the  furnace,  the  soft,  spongy 
reduced  iron  takes  up  carbon,  which  assists  greatly  in 
bringing  it  finally  into  the  molten  condition.  The  pig 
iron  formed  is  therefore  a  combination  of  iron  with  carbon, 
which  melts  much  more  readily  than  pure  iron. 

Blast  furnaces  were  formerly  open  at  the  top,  permit- 
ting the  waste  gases  and  products  of  the  combustion  and 
reduction  to  escape.  At  present,  as  the  illustration  shows, 
the  top  is  closed  and  the  gases  are  led  down  through  a 
large  pipe  from  which  they  are  conveyed  to  a  furnace  and 
burned  to  heat  the  air  blast  forced  in  through  the  tuyeres. 
This  exit  pipe  is  shown  at  p,  and  the  entrance  of  one  of 
the  tuyeres  at  m. 

Practically,  three  varieties  of  iron  are  recognized, 
viz.  :  Cast  iron,  wrought  iron  and  steel* 

Cast  or  Pig  Iron  is  characterized  by  containing  sev- 
eral per  cent  of  foreign  substances,  of  which  carbon  is 
essential  and  by  far  the  most  important.  Traces  of 
sulphur,  phosphorus,  silicon  and  manganese  are 
always  present.  The  first  two  in  more  than  very  minute 
traces  are  very  objectionable,  as  they  render  the  iron 
brittle.  Carbon  exists  in  cast  iron  in  two  forms  known  as 
combined  carbon  and  graphitic  carbon.  In  cast  iron  with 
a  gray  fracture  the  graphite  carbon  predominates.  When 
such  iron  is  dissolved  in  acids  the  carbon  is  left  in  the  free 
condition.  Cast  iron  in  which  the  carbon  is  in  chemical 
combination  has  a  white  fracture.  This  iron  dissolves  in 
acids  with  evolution  of  hydrocarbon  gases  having  a  pecu- 
liar characteristic  odor.  This  carbon  exists  in  the  iron  as 
carbide,  or  carbides,  as  there  are  doubtless  several  com- 
binations. These  may  be  broken  up  by  water  or  steam  at  a 
high  temperature  just  as  calcium  carbide  is  decomposed 
by  water  at  the  ordinary  temperature.  Cast  iron  is 
employed  practically  because  it  may  be  melted  at  a  rela- 
tively low  temperature  and  poured  into  molds. 

Wrought  Iron  is  usually  made  from  pig  iron  by  the 
operation  known  as  puddling.  The  pigs  are  melted  on  the 
hearth  of  a  reverberatory  furnace,  the  hearth  being  often 


380  GENERAL  CHEMISTRY. 

lined  with  oxide  of  iron.  The  oxygen  of  the  air  passing 
over  the  metal,  with  that  from  the  hearth,  gradually  burns 
out  the  carbon  to  form  CO,  while  the  silicon,  phosphorus 
and  sulphur  become  oxidized  and  with  a  little  of  the  iron 
form  a  slag.  These  operations  are  assisted  by  puddling  or 
stirring  the  melted  metal.  As  the  refining  progresses  the 
iron  becomes  thick  and  pasty;  it  is  then  worked  into  a  ball, 
withdrawn  from  the  furnace  and  hammered  and  rolled  to 
squeeze  out  the  slag  impurities.  To  further  improve  it,  it 
is  reheated  in  a  furnace  and  rerolled,  this  operation  being 
in  some  cases  repeated  several  times.  By  such  treatment 
the  metal  is  given  a  fibrous  structure,  greatly  increasing  its 
strength  and  tenacity.  Wrought  iron  is  as  nearly  pure  iron 
as  can  be  made.  It  often  contains  99.5  per  cent  of  real 
iron,  the  rest  being  made  up  of  traces  of  several  impurities, 
and  is  characterized  by  being  malleable,  ductile  and  easily 
welded.  It  requires  a  very  high  temperature  for  melting. 

Steel  is  chemically  a  product  between  wrought  iron 
and  cast  iron,  as  it  always  holds  a  certain  small  amount  of 
carbon  in  combination.  The  proportion  of  carbon  varies 
indifferent  kinds  of  steel  between  0.15  per  cent  and  1,5 
per  cent.  Several  methods  are  in  use  for  the  production 
of  steel.  In  one  process  bars  of  wrought  iron  are  heated  in 
contact  with  powdered  charcoal,  by  which  means  carbon  is 
absorbed  and  combined,  producing  what  is  known  as  blister 
steel,  from  the  appearance  of  the  bars  at  the  end  of  the 
heating.  This  is  melted  in  a  plumbago  crucible,  forming  a 
much  more  uniform  mass,  which  is  poured  into  molds  form- 
ing crucible  steel  or  cast  steel.  This  is  employed  for  many 
purposes.  At  the  present  time  the  largest  quantities  of 
steel  are  made  directly  from  pig  iron  by  the  Bessemer  process, 
or  by  the  Siemens  open  hearth  process. 

Bessemer  Steel  is  made  from  pig  metal  by  a  series  of 
comparatively  simple  operations.  The  pig  iron  is  melted 
in  a  furnace  and  run  into  a  vessel  called  a  converter, which 
is  furnished  with  a  perforated  bottom  and  tight  air  cham- 
ber below.  Through  the  perforations  air  is  blown  under 
considerable  pressure  into  th^  molten  metal.  It  speedily 
unites  with  and  burns  out  practically  all  of  the  carbon, 


GENERAL  CHEMISTRY.  381 

leaving  a  liquid  mass  corresponding  to  wrought  iron.  Into 
this  is  run  a  known  weight  of  melted  spiegeleisen,  which 
is  a  cast  iron  containing  manganese  and  very  rich  in  car- 
bon. The  proportion  of  carbon  in  this  having  been  deter- 
mined by  previous  analysis,  it  is  possible  to  take  enough 
of  it  to  produce  with  the  iron  in  the  converter  a  steel  of 
the  desired  composition.  The  manganese  leaves  the  iron 
and  goes  into  the  slag  always  formed  in  the  operation.  The 
converter  employed  is  usually  a  pear  shaped  vessel,  built 
of  iron  plates  and  lined  with  a  silicious  material  called 
ganister,  for  some  kinds  of  pig  metal,  or  with  a  mixture 
of  lime  and  magnesia  for  metal  rich  in  phosphorus.  It  is 
made  to  rotate  on  trunnions  so  that  it  may  be  tipped  to 
receive  the  charges,  and  again  at  the  end  of  the  operation 
to  pour  out  the  finished  steel.  As  much  as  fifteen  to  twenty 
tons  of  pig  iron  may  be  taken  as  a.  charge  and  converted 
into  steel  in  half  an  hour  in  one  of  these  converters.  Bes- 
semer steel  has  displaced  wrought  iron  tor  many  struc- 
tural purposes. 

Open  Hearth  Steel  is  made  by  melting  pig  iron  along 
with  iron  oxide  or  scrap  iron  in  the  hearth  of  a  kind  of 
reverberatory  furnace  resembling  a  puddling  furnace. 
The  pig  iron  becomes  decarbonized  to  the  right  extent  by 
properly  conducting  the  operation.  In  these  furnaces  a 
gaseous  fuel  must  be  used,  and  this  may  be  obtained  in  a 
simple  manner  from  coal,  or  from  other  source.  The  pro- 
duction of  steel  by  this  method  requires  more  time  than  is 
the  case  with  the  Bessemer  process,  but  the  quality  is  in 
general  more  uniform. 

Malleable  Iron  articles  are  made  by  casting  in  the 
usual  manner.  They  are  then  imbedded  in  powdered 
oxide  of  iron  and  heated  to  a  low  red  heat  about  two  days. 
In  this  manner  a  part  of  the  carbon  13  removed  from  the 
cast  iron  and  the  articles  become  soft  enough  and  malleable 
enough  to  be  worked  in  various  ways.  It  will  be  seen 
that  malleable  iron  is  a  cheap  substitute  for  wrought  iron. 

In  1896  the  world's  production  of  pig  iron  was 
34,110,814  tons,  of  which  the  United  States  furnished 
9,657,984  tons  and  Great  Britain  9,570,242  tons.  The 


382  GENERAL  CHEMISTRY. 

world's  production  of  steel,  in  1896,  was  19,339,244  tons,  of 
which  the  United  States  furnished  6,252,518  tons,  and 
Great  Britain  about  4,730,000  tons. 

Properties  and  Uses  of  Iron.  These  are  too  well 
known  to  call  for  extended  discussion  here.  Iron  is  very 
malleable  and  ductile  and  is  a  moderately  good  conductor 
of  electricity.  In  the  form  of  steel  it  takes  a  high  temper, 
that  is,  may  be  hardened  to  almost  any  desired  degree. 
Soft  iron  may  be  temporarily  magnetized  and  very 
strongly,  but  on  withdrawal  of  the  magnetic  influence  the 
property  is  lost.  Steel,  on  the  other  hand,  becomes  per- 
manently magnetized.  At  a  high  temperature  iron 
decomposes  steam,  liberating  hydrogen  and  leaving  the 
black  oxide,  Fe3O4.  When  heated  to  a  high  temperature 
and  plunged  into  a  jar  of  oxygen  it  combines,  forming  the 
same  oxide.  Iron  forms  two  classes  of  compounds,  ferrous 
and  ferric.  Its  power  of  combining  with  other  metals  to 
form  alloys  is  limited,  but  a  few  such  combinations  are 
known. 

Iron  Oxides.  Three  are  well  known.  Ferrous  oxide, 
FeO,  is  not  readily  prepared  or  kept  in  pure  condition.  It 
may  be  made  by  reduction  of  ferric  oxide  by  hydrogen  at 
a  temperature  of  300°.  When  exposed  to  air  or  oxygen  it 
oxidizes  immediately.  Ferric  oxide,  or  sesquioxide,  Fe2O3, 
occurs  in  nature  in  the  important  ores  mentioned.  It  is 
made  easily  by  heating  precipitated  ferric  hydroxide,  or  by 
the  calcination  of  dried  green  vitriol,  as  in  the  manufacture 
of  fuming  sulphuric  acid.  In  this  form  it  is  known  as  jew- 
elers' rouge  and  is  used  for  polishing.  Iron  is  trivalent  in 
this  and  other  ferric  compounds.  The  black  oxide,  magnetic 
oxide,  ferroso-ferric  oxide,  Fe3O4,  occurs  in  nature  as  an 
iron  ore  and  may  be  made  artificially.  It  may  be  looked 
uponas  a  combination  of  ferrous  andferric  oxide, FeO. Fe2O3. 
It  is  easily  made  by  passing  steam  over  red  hot  iron.  Such 
a  coating  is  often  formed  on  iron  to  keep  it  from  rusting,  as 
in  the  Bower-Barff  process. 

The  reduction  of  ferric  oxide  to  metallic  iron  is  illus- 
trated by  the  following  experiment: 


GENERAL  CHEMISTRY. 


Ex.  171.  Arrange  the  apparatus  as  shown  in  the  next  figure.  Gen- 
erate hydrogen  in  the  usual  manner,  and  dry  it  by  passing  it  through 
the  bottle  containing  strong  sulphuric  acid.  Charge  the  hard  glass  tube 
with  a  few  grams  of  ferric  oxide  (jewelers'  rouge).  The  outer  end  of  the 
tube  is  closed  with  a  perforated  stopper,  through  which  passes  a  narrow 
glass  tube  for  the  escape  of  the  hydrogen  and  steam.  After  the  appa- 
ratus is  arranged  allow  the  gas  to  stream  through  some  minutes  to  expel 
the  air  thoroughly.  Then  carefully  heat  the  tube  by  a  Bunsen  burner, 
and  keep  hot  until  steam  no  longer  escapes  from  the  open  end.  Mean- 
while a  good  current  of  hydrogen  must  continue  to  pass.  When  the 
reduction  is  complete  remove  the  lamp,  but  allow  the  product  to  cool 


FIG.    33. 


thoroughly  in  the  stream  of  hydrogen  before  taking  apart  the  apparatus. 
The  reduced  iron  is  left  as  a  fine,  dark  gray  powder,  which  burns  readily 
when  thrown  into  the  flame  of  a  Bunsen  burner. 

The  reduction  follows  according  to  the  equation: 

Fe203+3H2=3H20+2Fe. 

When  the  reduction  is  carried  out  at  a  proper  temperature,  a  low 
red  heat,  the  powder  possesses  the  property  of  burning  spontaneously 
when  thrown  into  the  air. 

Iron  Hydroxides.  Ferrous  hydroxide,  Fe(OH)2,  may 
be  made  as  a  greenish  precipitate  by  adding  an  alkali  solu- 
tion to  one  of  a  ferrous  salt.  In  absence  of  air  the  precip- 
itate is  almost  white,  but  it  oxidizes  quickly,  turning  green- 


384  GENERAL  CHEMISTRY. 

ish  brown.     Ferric  hydroxide,  Fe(OH)3,  is  best  made    by 
precipitating  a  ferric  salt  with  ammonia  water. 

FeCl3+3NH4OH^Fe(OH)3-f3NH4Cl. 

It  is  reddish  brown  and  soluble  in  strong  solutions  of 
ferric  salts.  Such  a  solution  exists  in  the  mixture  known 
as  dialyzed  iron.  A  crude  ferric  hydroxide  is  often 
employed  in  purifying  coal  gas  by  absorbing  sulphur  com- 
pounds. 

Iron  Sulphides.  Ferrous  sulphide,  FeS,  is  made  by  melt-  * 
ing  a  mixture  of  sulphur  and  scrap  iron  or  by  precipitat- 
ing a  ferrous  solution  by  means  of  ammonium  sulphide.  It 
is  a  black  substance  and  is  decomposed  by  acids  with  lib- 
eration of  hydrogen  sulphide.  Ferric  sulphide,  FeS2,  occurs 
as  pyrite  and  is  valuable  on  account  of  the  sulphur  it  con- 
tains. It  is  employed  largely  in  the  manufacture  of  sul- 
phuric acid.  The  native  substance  is  found  in  brassy  yel- 
low cubes  or  octahecfra. 

Iron  Chlorides.  Ferrous  chloride,  FeCl2,  may  be 
obtained  in  solution  by  dissolving  the  metal  in  hydro- 
chloric acid.  It  crystallizes  with  4H*2O  in  green,  very  solu- 
ble crystals.  The  pure  dry  substance  is  made  by  passing 
hydrochloric  acid  gas  over  hot  iron.  Ferric  chloride, 
FeCl3,  is  obtained  in  solution  by  dissolving  iron  in 
aqua  regia  and  repeatedly  evaporating  with  hydrochloric 
acid,  or  by  dissolving  in  hydrochloric  acid  and  passing  in 
chlorine.  It  is  very  soluble  in  water  and  also  in  alcohol 
and  ether.  The  anhydrous  substance  is  obtained  in  the 
form  of  crystalline  scales  by  passing  chlorine  gas  over  red 
hot  iron.  Ferric  chloride  is  one  of  the  most  important  of 
the  iron  salts. 

Ferrous  bromide,  FeBr2,  and  iodide,  FeI2,  are  well 
known  and  are  employed  in  medicines  in  syrups. 

The  reduction  of  a  ferric  to  a  ferrous  salt  and  the 
reverse  reaction  are  illustrated  by  the  following  experi- 
ments : 

Ex.  172.  Heat  a  few  cubic  centimeters  of  a  ferric  chloride  solu- 
tion in  a  test-tube  and  add  a  little  hydrochloric  acid  and  some  small 


GENERAL  CHEMISTRY.  385 

fragments  of  zinc.  A  lively  evolution  of  hydrogen  follows,  and  the  color 
of  ihe  solution  becomes  lighter.  In  a  short  time  it  disappears  almost 
completely,  leaving  a  ferrous  solution  with  a  faint  green  tinge.  The 
hydrogen  set  free  by  the  acid  and  zinc  acts  on  the  FeCl3,  removing  part 
of  its  chlorine. 

FeCl3-f-H=FeCl2+HCl. 

Ex.  173.  Dissolve  about  5  Gm.  of  green  vitriol  in  water,  in  a 
beaker,  and  add  3  or  4  Cc.  of  dilute  sulphuric  acid  (1  to  5)  and  about  10 
drops  of  strong  nitric  acid.  Heat  the  mixture  to  boiling  in  a  fume 
closet,  and  observe  the  change  of  color.  Red  fumes  escape,  and  when 
the  action  ceases  the  salt  described  below  is  left  in  solution. 

Iron  Sulphates.  Crystallized  ferrous  sulphate, 
FeSO4.7H2O,  is  a  common  substance,  known  as 
green  vitriol.  It  may  be  made  by  dissolving  iron  in 
dilute  sulphuric  acid  and  evaporating  to  crystalliza- 
tion or  by  the  oxidation  of  iron  sulphides  in  moist 
air.  The  salt  is  very  soluble  in  water,  but  not  in 
alcohol.  It  is  employed  in  the  preparation  of  many 
other  compounds  of  iron,  and  especially  in  the  pro- 
duction of  inks  and  pigments  used  in  dyeing.  A  crude  sul- 
phate is  used  as  a  deodorizer.  When  carefully  heated  to 
about  140°  it  loses  6H2O  and  leaves  a  colorless  salt.  The 
last  molecule  of  water  is  expelled  at  a  much  higher  heat. 
With  ammonium  sulphate  it  forms  the  important  com- 
pound (NH4)2SO4  FeSO4.6H2O.  This  is  much  more 
stable  than  green  vitriol.  Ferric  sulphate,  Fe2(SO4)3,  is 
obtained  in  solution  by  oxidizing  ferrous  sulphate  by  aid 
of  nitric  acid  after  addition  of  more  sulphuric  acid.  This 
solution  is  used  in  pharmacy.  Ferric  sulphate  crystallizes 
with  ammonium  sulphate  in  the  proper  proportion,  yield- 
ing/^r^0/«w,(NH4)2SO4.Fe2(SO4)3.24H2O.  This  salt  is 
stable  in  the  air.  A  similar  compound  is  formed  with 
potassium  sulphate. 

Iron  Nitrates.  The  ferrous  and  ferric  salts  are  known. 
Ferric  nitrate,  Fe(NO3)3,  is  employed  in  making  Prussian 
blue  for  dyeing. 


Iron  Phosphates. 
is  formed  as  a  light  colored  precipitate   by  precipitating  a 


386  GENERAL  CHEMISTRY. 

solution  of  ferrous  sulphate  with  sodium  phosphate.     It  is 
used  in  medicine.      Ferric  phosphates  are  also  known. 

Other  Compounds.  Iron  is  found  in  the  important 
substances,  potassium  ferrocyanide,  K4Fe(CN)6,  and  po- 
tassium ferricyanide,  K3Fe(CN)6.  These  may  be  looked 
upon  as  double  cyanides,  or  as  salts  of  ferrocyanhydric 
acid  and  ferricyanhydric  acid.  They  are  used  as  valuable 
laboratory  reagents  and  in  the  manufacture  of  Prussian 
blue  and  Turnbull's  blue. 

Recognition.  The  hydrated  ferrous  salts  are  usually 
light  green.  The  ferric  salts  yield  yellow  to  brown  solu- 
tions. Ferrous  salts  give  with  ammonia  light  greenish 
ferrous  hydroxide,  while  ferric  salts  give  a  precipitate  of 
reddish  ferric  hydroxide.  With  potassium  ferricyanide 
ferrous  salts  give  a  deep  blue,  while  ferric  salts  yield  a 
deep  blue  precipitate  with  potassium  ferrocyanide. 

NICKEL. 

Occurrence.  This  metal  occurs  usually  as  sulphide 
associated  with  other  metallic  sulphides.  Nickel  blende 
NiS,  nickel  glance,  Ni(AsS)2  and  kupfer-nickel,  NiAs, 
and  a  complex  nickel-magnesium  silicate  are  the  most  im- 
portant ores. 

History.  Nickel  has  been  known  since  1751,  when  it 
was  discovered  among  German  ores  worked  for  copper. 
After  that  time  the  metal  was  found  rather  widely  distrib- 
uted and  is  now  produced  from  mines  in  Germany,  New 
Caledonia  and  Canada. 

Metallurgy.  The  processes  employed  in  the  prepara- 
tion of  pure  nickel  from  its  ores  are  in  the  main  complex 
because  of  the  peculiar  nature  of  the  native  compounds. 
In  the  simplest  case  the  sulphide  is  roasted  to  form  oxide 
and  this  is  reduced  with  carbon  in  the  usual  manner.  But 
as  a  rule  much  more  difficult  processes  are  necessary,  and  a 
description  of  them  lies  beyond  the  scope  of  this  book. 


GENERAL  CHEMISTRY.  387 

No  accurate  figures  can  be  given  of  the  production  of 
pure  nickel  at  the  present  time.  In  1896  this  country 
furnished  about  1,800  tons,  chiefly  from  Canadian  ores. 
The  world's  total  production  was  about  four  times  that 
amount. 

Properties.  Nickel  is  a  very  hard  metal  which  takes 
a  high  polish  and  corrodes  in  the  air  only  superficially. 
It  is  but  slowly  dissolved  by  acids,  and  alkalies  may  be 
fused  in  contact  with  it  without  danger  of  much  corro- 
sion. By  special  treatment  the  metal  has  been  made  mal- 
leable and  ductile,  and  in  this  form  has  found  many  uses. 
It  alloys  with  several  metals  readily. 

Uses.  It  is  employed  mainly  as  a  coating  for  iron  in 
the  form  of  nickel  plate.  It  is  also  used  in  making  sev- 
eral important  alloys.  German  silver  contains  copper,  zinc 
and  nickel.  The  5  cent  pieces  of  this  country  contain  75 
parts  of  copper  and  25  of  nickel.  A  combination  with  iron 
known  as  nickel-steel  is  also  important. 

Compounds  of  Nickel.  The  sulphate,  NiSO4.7H2O, 
and  the  ammonium-nickel  sulphate,  (NH4)2  Ni(SO4)2. 6H2O, 
are  the  most  important  compounds.  The  latter  is  com- 
monly used  in  nickel  plating.  Hydrated  nickel  salts  are 
green  to  greenish  blue,  while  the  anhydrous  salts  are  yel- 
low. Many  other  nickel  compounds  are  known,  but  they 
have  no  technical  importance. 

Recognition.  When  nickel  compounds  are  fused  in  a 
borax  bead  this  becomes  yellowish  brown.  In  alkaline  so- 
lution nickel  compounds  yield  with  hydrogen  sulphide  a 
black  precipitate  of  NiS. 


COBALT. 

Occurrence.  This  metal  is  found  often  associated  with 
nickel  in  arsenic  or  sulphur  combination.  Free  cobalt  and 
nickel  are  found  in  meteorites.  An  arsenide  containing 
nickel,  cobalt  and  iron,  called  tin  white  cobalt,  is  the  most 
important  ore. 


388  GENERAL  CHEMISTRY. 

Metallurgy.  The  separation  of  cobalt  from  its  ores 
is  difficult  and  the  industry  is  not  an  important  one. 

Properties  and  Uses.  Cobalt  is  a  white  metal  with 
a  reddish  tinge.  It  is  magnetic  and  very  hard  and  in  the 
form  of  wire  possesses  great  strength.  It  is  possible  to 
deposit  cobalt  as  a  coating  on  other  metals  as  nickel  is  de- 
posited, but  this  industry  has  not  yet  become  important. 

Cobalt  Compounds.  The  hydrated  salts  are  rose  red> 
but  when  deprived  of  their  water  of  crystallization  they 
are  deep  blue.  The  monoxide,  CoO,  is  a  brownish  black 
powder.  The  best  known  salts  are  the  sulphate, 
CoSO47H2O,  the  chloride,  CoCl26H2O,  and  the  nitrate, 
Co(NO3)2.6H2O.  The  crude  silicate,  known  as  smalt,  is 
largely  employed  in  coloring  glass  and  pottery  deep  blue. 

Cobalt  forms  a  large  number  of  very  complex  com- 
pounds with  ammonia  which  are  known  as  cobalt  amins. 
These  combine  with  acids  forming  salts  which  are  mostly 
highly  colored.  With  a  solution  of  potassium  nitrite  in 
presence  of  acetic  acid  cobalt  salts  form  a  beautiful  crys- 
talline yellow  precipitate  of  potassium  cobalt  nitrite, 

K8Co(Nba)6. 

The  change  of  color  on  dehydration  of  a  cobalt  salt  is 
shown  in  the  next  experiment.  A  common  sympathetic 
ink  is  a  dilute  solution  of  the  chloride  or  nitrate  of  cobalt, 
the  use  of  which  is  here  illustrated. 

Ex.  174.  Dip  a  clean  pen  into  a  dilute  solution  of  cobalt  chloride, 
CoCl2.6H2O,  and  write  with  it  on  white  paper.  When  dry  the  writing 
is  almost  invisible.  Now  warm  the  paper  gently  and  carefully  to  avoid 
scorching  it.  Soon  the  writing  becomes  blue,  and  distinctly  visible.  If 
the  paper  is  allowed  to  stand  in  the  air  some  time  the  anhydrous  salt, 
CoCl2,  absorbs  moisture  and  returns  to  the  almost  colorless  form. 

Recognition.  Cobalt  compounds  are  recognized  by  the 
yellow  nitrite  precipitate  just  referred  to  and  by  forming  a 
deep  blue  bead  when  fused  with  borax.  This  last  behavior 
is  extremely  characteristic. 


CHAPTER  XXIII. 


THE  PLATINUM  GROUP  OF  METALS  :     RUTHENIUM, 

RHODIUM,  PALLADIUM,  OSMIUM,  IRIDIUM 

AND   PLATINUM. 


GENERAL  CHARACTERISTICS. 

THESE  six  metals  are  classed  together  in  the  eighth 
group  or  family  of  the  Periodic  arrangement,  forming 
the  last  part  of  the  fifth  and  eighth  periods.  The  group  of 
six  naturally  falls  into  two  subgroups  of  three  each,  as 
the  first  three  elements  have  atomic  weights  near  100  and 
the  last  three  weights  near  200.  The  proper  scientific 
classification  of  these  metals  offers  some  difficulties 
because  of  the  closeness  in  the  atomic  weights  in  each 
group,  but  for  practical  reasons  they  are  always  treated 
together. 

The  physical  constants  of   the  metals  are  shown  in  the 
following  table  : 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY. 

ATOMIC 
VOLUME. 

MELTING 
POINT. 

Ruthenium   .  . 

101  68 

12  20 

8  29 

1800° 

Rhodium  

103  01 

11  10 

8  51 

2000° 

Palladium  

106.36 

11.45 

9  29 

1500° 

Osmium  

190.99 

22  48 

8  49 

2500° 

Ir'dium 

193  12 

22  42 

8  61 

2200° 

Platinum 

194  89 

21  50 

9  06 

17750 

390  GENERAL  CHEMISTRY. 

PLATINUM. 

Occurrence.  The  six  metals  of  the  group  are  found 
associated  in  'the  ore  known  as  native  platinum  which 
occurs  in  granular  form  in  sandy  or  alluvial  deposits  in 
several  parts  of  the  world,  more  particularly  in  western 
Siberia.  Smaller  amounts  are  found  in  Australia,  in  Mexico, 
in  South  America  and  in  Borneo.  Platinum  is  by  far  the 
most  abundant  of  the  metals  in  this  ore,  making  up  75 
per  cent  of  it  usually.  The  occurrence  of  platinum  in 
other  ores  is  quite  unimportant. 

History.  This  ore  has  been  known  in  a  general  way 
since  the  middle  of  last  century,  but  it  assumed  no  techni- 
cal importance  until  the  English  chemist,  Wollaston,  dis- 
covered a  method  of  separating  the  platinum  and  bringing 
it  into  malleable  condition.  These  discoveries  were  util- 
ized by  the  London  firm  of  Johnson,  Matthey  &  Co.,  who, 
about  1810,  began  the  manufacture  of  platinum  articles  for 
technical  purposes.  Many  of  the  advances  in  the  art  of 
working  the  metal  have  been  made  by  this  firm. 

Metallurgy.  The  ore  is  generally  separated  from  the 
sands  containing  it  by  a  systematic  washing  process,  such 
as  was  referred  to  under  gold.  It  is  then  worked  up  by  one 
of  two  essentially  different  processes.  The  first  of  these  is 
a  modification  of  that  originally  devised  by  Wollaston  and 
consists  in  dissolving  out  the  metals  by  aqua  regia,  leaving 
sand  and  certain  insoluble  bodies  behind.  The  strong 
solution  of  the  chlorides,  after  evaporation  to  remove 
excess  of  acid,  and  dilution,  is  precipitated  with  an  excess 
of  ammonium  chloride,  which  throws  down  the  platinum 
in  the  form  of  a  crystalline  double  salt,  PtCl4.2NH4Cl, 
with  a  small  amount  of  the  analogous  iridium  salt.  The 
mother  liquor  contains  nearly  the  whole  of  the  other  metals 
and  a  little  of  the  platinum.  All  these  may  be  precipitated 
in  spongy  form  by  metallic  iron.  The  precipitate  so 
obtained  is  heated  with  hydrochloric  acid  to  free  from 
traces  of  iron,  dissolved  in  aqua  regia  and  then  reprecipi- 
tated  by  an  excess  of  ammonium  chloride  as  before.  The 
resulting  mother  liquor  is  used  for  the  preparation  of  the 


GENERAL  CHEMISTRY.  391 

rarer  elements,  while  the  mixed  platinum -iridium  precipi- 
tates are  washed  and  converted  into  metal.  This  is  done 
by  heat,  as  the  double  salts  are  easily  decomposed,  leav- 
ing a  spongy  mass,  which  may  be  brought  into  compact 
form  by  hammering  or,  commonly,  by  fusing  with  the 
oxyhydrogen  blowpipe  in  a  massive  stone  crucible.  The 
metal  obtained  is  not  pure  platinum  but  an  alloy  contain- 
ing a  little  iridium,  which  for  most  purposes  is  practically 
better  than  the  pure  metal.  The  manufacture  of  pure  plati- 
num is  an  operation  of  considerable  difficulty  and  will  not 
be  described  here. 

By  another  process, devised  by  Deville,in  France,  plati- 
num is  separated  from  the  ore  by  a  fusion  method.  The 
separation  is  effected  by  smelting  with  lead  sulphide  which 
combines  readily  with  platinum  but  takes  up  but  little  of 
the  osmium  and  iridium.  The  regulus  of  lead  and  plati- 
num obtained  is  cupelled  for  separation  of  the  lead  by  oxi- 
dation. The  resulting  platinum  is  sometimes  pure  enough 
for  use  directly  but  usually  is  subjected  to  further  refining 
operations. 

In  1896  the  total  production  of  platinum  was  about 
10,000  pounds,  and  chiefly  from  Ural  ores. 

Properties.  Platinum  is  very  ductile  and  malleable, 
is  a  good  conductor  of  heat  and  electricity  and  does  not 
oxidize  at  any  temperature.  It  does  not  dissolve  in  hy- 
drochloric or  nitric  acid,  but  in  a  mixture  of  the  two,  or  in 
liquids  which  liberate  chlorine.  When  alloyed  with  silver 
small  amounts  are  soluble  in  nitric  acid.  At  a  high  tem- 
perature platinum  is  attacked  by  fused  alkalies  and  phos- 
phorus, also  by  many  metals  and  heavy  metallic  salts. 
There  is,  therefore,  a  limit  to  the  number  of  substances 
which  may  be  heated  or  fused  in  platinum  vessels.  It  al- 
loys easily  with  many  of  the  heavy  metals. 

By  several  methods  platinum  may  be  brought  into  a 
finely  divided  form  known  as  platinum  black  or  platinum 
sponge,  and  in  this  condition  it  absorbs  gases  readily. 
Several,  applications  are  made  of  this  property  in  the  arts. 

Uses.     Platinum  is  largely  used  in  the  manufacture  of 


393  GENERAL  CHEMISTRY. 

laboratory  ware  for  a  great  variety  of  purposes,  also  in  the 
construction  of  large  stills  for  the  distillation  of  sulphuric 
acid  or  for  its  concentration  without  distillation.  These 
large  platinum  vessels  are  rendered  more  durable  by  being 
plated  with  gold.  Until  1845  the  metal  was  used  for  coin- 
age in  Russia. 

Platinum-Iridium  alloys  have  come  into  use  for  the 
production  of  several  articles.  Much  time  has  been  spent 
in  devising  methods  of  separating  both  metals  in  perfectly 
pure  condition,  and  alloys  may  now  be  obtained  containing 
them  in  any  desired  proportion.  A  combination  of  90 
parts  of  platinum  with  10  parts  of  iridium  is  very  hard 
and  withstands  a  high  temperature.  Because  of  these  and 
other  important  properties  it  is  used  in  the  manufacture  of 
the  standard  weights  and  measures  in  Great  Britain.  An 
alloy  with  15  parts  of  iridium  is  used  in  making  standard 
rules.  Most  of  our  laboratory  platinum  ware  contains 
iridium. 

Compounds  of  Platinum.  Many  of  these  are  known, 
but  few  have  any  practical  applications.  Platinous  chlo- 
ride, PtCl2,  is  known  and  a  combination  of  this  with  potas- 
sium chloride,  PtCl2.2KCl,  is  employed  in  photography 
under  the  name  of  potassium  chlorplatinite  to  give  the 
dark  or  black  effects  in  the  toning  of  prints.  The  most 
important  salt  of  platinum  is  the  tetrachloride,  PtCl4, 
which  is  obtained  by  solution  of  the  metal  in  aqua  regia. 
It  is  used  in  chemical  analysis  for  the  separation  of  potas- 
sium, because  under  certain  conditions  it  forms  the  insolu- 
ble compound,  PtCl4.2KCl.  The  corresponding  ammo- 
nium salt  has  been  referred  to  above,  as  made  in  the 
process  of  separating  platinum  from  its  ore. 

Recognition.  The  metal  is  usually  recognized  by  the 
formation  of  these  yellow  precipitates. 

The  Other  Platinum  Metals. 

Nothing  more  than  brief  mention  need  be  made  of 
these.  Iridium  is  chiefly  interesting  and  important  be- 


GENERAL  CHEMISTRY.  393 

cause  of  the  alloys  it  forms  with  platinum.  To  secure 
hardness  the  points  of  gold  pens  are  often  made  of  this 
metal.  Osmium  is  always  found  in  platinum  ores  and  also 
in  an  ore  with  iridium  alone.  It  has  been  obtained  in  com- 
pact form  and  is  the  heaviest  metal  known.  It  forms  a 
number  of  oxides,  two  of  which  have  acid  properties. 
Osmic  tetroxide,  OsO4,  or  perosmic  acid,  is  a  volatile 
substance  which  is  intensely  poisonous,  and  which  is  used 
as  a  stain  in  microscopic  work.  Ruthenium  and  rhodium 
compounds  have  no  practical  applications.  Palladium  is 
interesting  and  important  because  of  its  property  of  absorb- 
ing large  quantities  of  hydrogen,  of  which  it  will  take  up 
several  hundred  volumes  under  certain  conditions.  This 
property  is  applied  in  the  analysis  of  gases  containing 
hydrogen.  By  some  chemists  the  product  is  considered 
as  an  alloy  of  palladium  and  hydrogen;  others  conclude  that 
a  compound  of  the  formula  PdH  or  Pd3H2  may  be  formed, 
and  finally  it  may  be  looked  upon  as  merely  a  mechanical 
mixture. 


INDEX. 


Absorption  of  heat 59 

spectra 326 

Acetates 228 

Acetic  acid 228 

Acetylene 222 

Acid,  arsenic 199 

arsenous 197 

boric 178 

chloric 93,  143 

chlorous 92,  143 

fermentation 228 

fluosilicic 178 

hydrobromic 97 

hydrochloric 72 

hydrocyanic 229 

hydrofluoric 102 

hydriodic 100 

hypobromous. 98 

hypochlorous 143 

hypophosphorous 188 

hyposulphurous 155 

nitric 125  to  130 

nitrous 123 

perchloric 92,  143 

phosphoric 189 

producer    37 

prussic 229 

pyroarsenic 199 

pyroligneous 228 

selenic 170 

selenous 170 

silicic 173 

sulphuric 156 


Acid,  sulphurous 152,  154 

telluric 170 

tellurous 170 

Acids 256 

nature  of 104 

of  chlorine 90 

Air,  ammonia  in 119 

analysis 110,  111 

carbon  dioxide  in 118 

moisture  in 118 

ozone  in 119 

tests 116 

Albertus  Magnus 194 

Alcohol 227 

boiling  point 19 

Alcoholic  solutions 12 

Alkali-earth  group 310 

Alkalies 105 

Alkali  metals 271 

Alloys 256 

Alum 10,  339 

solubility  of 6 

Aluminum 337 

bronze 291,338 

chloride 339 

hydroxides 339 

oxide 339 

silicate 340 

sulphate 339 

Alums 340 

Amalgams 255 

Amalgamation  process. ...  295,  305 
Ammonia..  132 


396 


INDEX. 


Ammonia,  anhydrous 137 

caustic 135 

for  refrigeration   137 

in  air 119 

solubility  of 136 

Ammonium 132,  286 

carbonate 286 

chloride 3,  26,  134 

compounds 271 

hydroxide 1 35 

molybdate 191,  367 

nitrate 120,   134 

sulphate 134 

sulphide 167,  287 

Amorphous  carbon 204 

phosphorus. 186 

Anaesthesia 121 

Analysis  of  air 110,  111 

Analytical  reactions 22 

Anhydrous  ammonia 137 

Antidotes  for  arsenic 200 

Antimonic  acid 357 

Antimony 356 

and   hydrogen 358 

chlorides 358 

oxides 357 

sulphides 358 

Aquamarin 311 

Aqua  regia 131 

Argon 117 

Arsenates 199,  265 

Arsenic ...183,  193 

acid 199 

and  copper 198 

and  hydrogen 194 

and  sulphur 200 

and  the  halogens 199 

oxide 198 

poisoning,  antidotes 200 

tests..  ...   196 


Arsenical  pyrite 194 

Arsenites 198,  265 

Arsenous  oxide 197 

Arsine 194 

Artificial  ice 137 

Ash  of  seaweed 7 99 

Atmosphere 106,  109 

Atmospheric  ozone 43 

Atomic  combinations 243 

theory 78 

volumes 253 

weights 79,  231 

weights,  table 80,     81 

Atoms 77 

Avogadro 232 

Avogadro's  law 241 

Azoimid 138 

Azote 106 

Baking  powder 27 

soda 27 

Balance  in  chemistry 37 

Balard 94 

Barium 318 

chlorate 92,  320 

chloride 319 

dioxide 64 

hydroxide „ ,   319 

oxide 319 

sulphate 319 

sulphite 151 

Bases 105,  2o6 

Bauxite 339 

Beryl 311 

Beryllium 311 

Berzelius '. . .  .79,  233 

Bessemer  steel 380 

Bismuth 359 

and  the  halogens 360 

nitrates. .  360 


INDEX. 


397 


Bismuth  oxides 360 

sulphides 361 

Black 270 

Black  lead 203 

Blast  furnace 377 

Bleaching  agent 68 

by  sulphurous  oxide 150 

powder 76,     91 

Blende 328 

Blue  vitriol 10,  292 

Bohemian  glass 176 

Boiler  scale 210 

Boiling  point 60 

of  alcohol 19 

of  ether 19 

Boneblack  filters 206 

Bottle  glass 176 

Borates 180,  263 

Borax 11,   179,   280 

bead 179 

Borocalcite 179 

Boron 172,  178,  336 

and  oxygen 178 

detection  of 182 

nitride 182 

Brand 183 

Brass 291 

Braunite 370 

Brimstone 147 

Bromides 94 

Bromic  acid 95 

Bromination  process  306 

Bromine <)0,     93 

constants 96 

Bronze 291 

Bunsen 270,  319 

Cadmium 331 

salts 331 

Caisium. .  .   286 


Calamine 328 

Calcium 314 

bicarbonate 209 

carbide 222,  317 

carbonate 316 

chlorate 285 

chloride 316 

hydroxide 315 

hypochlorite 91 

light 40 

oxalate 318 

oxide 315 

phosphate 184 

sulphate 317 

tartrate 11 

Calculations  from  equations. ..     88 

Calico,  bleaching  of 68 

Calomel  334 

Calorimeter 268 

Camphor 2 

Cannizzaro 244 

Capacity  for  heat 59 

Carbon 202 

and  chlorine 229 

and  hydrogen 220 

and  nitrogen 229 

and  oxygen   208 

and  sulphur 230 

Carbonates 209,  212,  259 

Carbon  dioxide 208 

dioxide  in  air ...118,  211 

disulphide 230 

group 344 

hydrogen  and  oxygen 223 

monoxide 212 

Carbonic  acid 210 

Carnallite 281 

Cast  iron 379 

Cassiterite  3l5 

Caustic  ammonia 135 


398 


INDEX. 


Caustic  soda 

Carat 

Cavendish 81,  44,  51, 

Cave  niter 131, 

Celestine 

Celsius  scale 

Centigrade  scale 

Cerium 

Chalk,  precipitation  of 

Chamber  acid 

Change,  chemical  and  physical. 
Change  of  state 

by  solution 

Chaptal 

Charcoal 

and  niter 

in  niters 

Chemical  changes 20, 

problems 88, 

Chili  saltpeter 98,< 

Chloric  acid   91, 

Chlorine 

and  carbon. 

and  hydrocarbons  

and  metals 

and  oxygen 

and  sulphur 

as  bleaching  agent 

compounds  

constants 

dioxide 

monoxide 

preparation G6, 

solubility  of 

tests  for 

trioxide 

uses 

water 

Chloroform 

Chlorous  acid. . 


275  Chlorination  process 306 

308  Chromates 265,  365 

125  Chrome  green 363 

317  orange 366 

318  yellow 366 

62  Chromium 363 

62  alum 8,  364 

352  chloride 364 

14  group 362 

160  oxide 109 

20  oxides 363 

2  Chrysoberyl 311 

21  Cinnabar 332 

106  Clarke 79 

204  Classes  of  metallic  compounds.   256 

128  Classification  of  elements 249 

206  Clay 337 

25  Coal  gas 134 

89  Cobalt 387 

280  compounds 388 

92  Coefficient  of  expansion. 112 

66  Coke 204 

229  Collecting  gases 32 

69  Colored  fires 285 

68  Coloring  glass 177 

90  Columbium 355 

169  Combination  of  gas  volumes. ..  232 

68  Combined  acid 164 

90  Combining  weights 140 

72  Combustion  in  chlorine 68 

91  in  nitric  acid 127 

90  in  oxygen 35 

67  of  candle 69 

70  of  gases 28 

71  spontaneous 36 

91  Composition  of  \vater 51 

72  Compounds  of  chlorine 90 

70  Compounds  of  nitrogen 120 

229  Condensation  of  chlorine 72 

92  of  hydrochloric  acid 75 


INDEX. 


899 


Condenser,  Liebig's 18 

worm 18 

Courtois 98 

Conditions  of  chemical  change.     25 

of  oxidation .  .    38 

Correction  for  pressure 115 

for  temperature 112 

Corrosive  sublimate 334 

Copper 288 

and  iron  scale 20 

and  nitric  acid 122 

arsenite 198,  294 

chlorides 292 

group 238 

hydroxides 292 

mines 289 

oxide 55 

oxides 291 

plating 293 

sulphate 292 

and  crystallization ....        8 

in  Deacon  process 71 

turnings 149 

Crown  glass 176 

Crude  potash 283 

Cruikshank 41 

Cryolite 101,  337,  341 

Crystallization 7 

fractional 9 

of  alum 8 

of  chromtfalum 8 

of  copper  sulphate 8 

water  of 10 

Crystals 8 

Cyanogen 229 

Cyanide  process 306 

Dalton 77 

Dalton's  weights 231 

Davy 66,219,272,  319 


Deacon  process 71 

Decomposition  of  water .  48 

Deflagrating  spoon 35 

Deh}dration  of  crystals 11 

Dephlogisticated  air 38 

Determination  of  atomic  weight  247 

of  specific  heat 268 

Developers 303 

Dextrose 224 

Diamond 202 

Dichromates 266 

Diffusion  of  hydrogen 49 

Dioxide,  hydrogen 63 

Direct  vision  spectroscopes. . . .  327 

Disilicates 267 

Distillates 17 

Distillation 16 

fractional 19 

Distilling  apparatus 17,  18,  19 

Dithionic  acid 164 

Double  decomposition 87 

fluorides 101 

Drummond  light 40,  219 

Dry  reactions 26 

Dulong  and  Petit 237 

Electrolysis 52 

Electrolytic  equivalents 239 

Electroplating 293 

Elements  and  compounds 76 

Emerald 311 

green 198 

Emery 339 

Epsom  salts 312 

Equations 82 

Etching  glass ...  102 

Ether,  boiling  point 19 

Ethylene 221 

Eudiometer 58 

corrections 114,  115 


400 


INDEX. 


Expansion  of  gases 112 

Experiment 2 

Experimental  science 1 

Faraday 239 

Fehling  test 225 

Felspar 337 

Fermentation 225 

Ferric  alum 385 

chloride 384 

oxide ; 382 

sulphate '. 385 

Ferrous  bromide 384 

chloride 384 

sulphate 12,   86,  162,  385 

sulphide 165,  384 

Filters 5 

Filtration 5,  63 

Fire  damp 221 

Fireworks 285 

Flint  glass 176 

Flowers  of  sulphur 147 

Fluorides 101 

Fluorine 90,  101 

and  glass 102 

Fluorspar 101 

as  flux 103 

Fluosilicic  acid 178 

Formulas 82 

graphic 143 

results  of  experiment 83 

Fractional  crystallization 9 

distillation  ." 19 

Franklinite 329 

Freezing  point 59 

Fuel  gas 218 

Fuming  nitric  acid 127 

Fusible  metals 256 

alloys 360 


Galena 348 

Gallium 342 

Galvanized  iron 330 

Gas  burners 218 

Gases,  collection  of 32 

in  atmosphere 117 

Gasholder 34 

manufacture 216 

problems 106 

reactions 27 

volume  combinations. . . . . .  232 

volume  reduction 112 

Gay  Lussac 51 ,  232 

tower 159 

Generalities 103 

General  reactions 24,  25 

Germanium 345 

Glass 176 

negatives 302 

soluble 175 

Glauber 125 

Glauber's  salt • 10,  280 

Glazes 341 

Glover  tower 159 

Glucinum 311 

Glucose 225 

Gold 305 

alloys 308 

chloride 309 

refining 307 

Gram 62 

Graphic  formulas 143 

Graphite 203 

Green  vitriol 12,  385 

Gum  dextrin 12' 

Gunpowder 128,  284 

Gypsum 317 

Hard  water..                         ..63,  210 


INDEX. 


401 


Halides 259 

Halogens  and  phosphorus 191 

Heat  absorption 59 

capacity 59 

of  combustion 39 

unit 59 

units 39 

Heavy  spar 319 

Helium 117 

Humboldt 51 

Hydrazin 138 

Hydrocarbons 220 

Hydrochloric  acid 66,  72 

constants 74 

Hydrofluoric  acid 102 

Hydrogen 30,  43 

and  nitrogen 132 

and  oxygen 55,  56,  57 

arsenide 194 

constants 50 

diffusion  of 49 

dioxide 03 

preparation 64 

reducing  power 50 

sulphide 165 

Hydronitric  acid 138 

Hydroquinon 303 

Hydroxides 2~>8 

Hydroxylamin 137 

Hypo  acids 144 

I  lypochlorites 90 

Hypophosphites 188 

Ice,  artificial 137 

Illuminating  gas 214 

Incandescent  lamps ....  353 

Indium 342 

Induction  coil 58 

Inflammable  air 44 

lodic  acid..                                   .  101 


Iodine 2,  90,  98 

and  hydrogen 100 

pentoxide 101 

Iridium 392 

Iron 376 

and  oxygen 36 

chlorides 304 

group 375 

hydroxides 383 

nitrates 385 

ores 376 

oxides 382 

phosphates 385 

sulphates 385 

sulphides 384 

Isomorphism 237 

Isomorphous  substances 9 

Kaolin 340 

Kelp 98 

Kindling  temperature 39 

King's  yellow 200 

Lampblack 207 

Lanthanum 343 

Laplace 270 

Latent  heat 59 

Laughing  gas 121 

Lavoisier 106,  125.  202,  270 

Law  of  Avogadro 241 

Lead 3J8 

acetate 351 

bromide 351 

carbonate  351 

chloride 351 

chromate 351 

iodide 351 

nitrate 124,  351 

oxides 350 

sulphate 350 


402 


INDEX. 


Lead  sulphide 350 

tree 350 

Leaden  pans 1 59 

Leblanc  process 277 

Liebig  condenser 18 

Lime 315 

Limelight 40,  219 

Liquid  carbon  dioxide 212 

Litharge 250 

Lithium .  ..  272 

in  water 325 

Lunar  caustic 300 

Magnesia  alba 312 

Magnesium 312 

carbonate 313 

chloride 313 

light 312 

oxide 313 

phosphate 314 

sulphate 14,  314 

Malleable  iron 381 

Manganates 266,  373 

Manganese 369 

chlorides 371 

dioxide 32,  87 

hydroxides 371 

oxides 370 

sulphates 372 

Marble  dust 4 

Marsh  gas 220 

test 196 

Matches 185 

Matthiessen 272 

Measuring  gases 58 

Mendelejeff 251 

Mercuric  chloride 26,  334 

iodide 26,  334 

oxide 31,  333 

nitrate. .                                 .  335 


Mercuric  sulphate 334 

sulphide   H33 

Mercurous  chloride 334 

iodide 334 

oxide 333 

Mercury 332 

Metallic  elements 255 

properties. . .    : 255 

Metallurgy  of  copper 290 

of  silver 295 

Metasilicates 267 

Methane 220 

Methyl  orange 133 

Meyer 251 

Michigan  salt  wells 94 

Milk  of  lime 315 

Mitscherlich 234 

Moissan 101 

Moisture  in  air 118 

Molecular  weight 82 

Molecules 77.  82 

Molybdenum 366 

Monazite 352 

Mortar 316 

Mother  liquor 9 

Mother  of  vinegar 228 

Multiple  proportions 141 

Natural  arrangement ...   251 

Natural  waters 62 

Nature  of  acids 104 

of  chemical  changes 22 

of  solutions 4 

Native  platinum 390 

Negative  elements 249 

Negatives 302 

Nickel.- 386 

compounds 387 

Niter,  solubility  of 6 

Nitrates..  ..130,  262 


INDEX. 


403 


Nitrates  in  soils 131 

reduction  of 124 

Nitric  acid 124,  125 

and  copper 122 

and  starch 123 

combustion  in 127 

fuming 127 

uses 130 

Nitrides 139 

Nitrites 262 

Nitrogen 106 

and  halogens 139 

and  hydrogen 132 

and  oxygen 120,  141 

compounds  of 120 

chloride 139 

dioxide   122 

group 354 

iodide 139 

monoxide 120 

pentoxide 124 

separation  of 107 

tetraoxide 124 

trioxide '.....  123 

Nitroso-sulphuric  acid 157 

Nitrosyl  chloride 131 

Nitrous  acid 123 

oxide 121 

Nitroxyl  chloride 131 

Nomenclature  of  acids  and  salts  145 

Nonmetallic  elements 249 

Normal  reaction 129 

Oil  of  vitriol 162 

Open  hearth  steel 381 

Orthoclase 174 

Orthophosphoric  acid 190 

Orthosilicates 267 

Osmium  . .  .393 


Oxalic  acid 213 

solubility  of 6 

Oxidation  by  saltpeter 128 

conditions  of 38 

in  soils 131 

of  magnesium 21 

of  metals 20 

of  sulphites 154 

Oxide  of  copper 55 

of  sulphur 35 

Oxides 257 

of  carbon 48 

of  chlorine 90 

of  nitrogen 141 

Oxygen 30 

and  arsenic 197 

and  boron 178 

and  chlorine 90 

and  nitrogen  120 

and  silicon 172 

and  sulphur 149 

constants 40 

salts  of  halogens 262 

uses 40 

Ozone 40 

generators 43 

in  air ..43,  119 

paper 42 

Palladium 393 

Paraffin  bottles " 103 

Paris  green 198 

Parkes1  process 296 

Pearl  ash 283 

Pentathionic  acid 164 

Perchlorates 93 

Perchloric  acid 92 

Periodic  arrangement 251 

functions. . .                            .  250 


404 


LVD  EX. 


Peroxide  of  hydrogen 63 

Persulphuric  acid 164 

Permanganates 372 

Petroleum 220 

Phenol-phthalein 133 

Phlogiston 37 

Phosphates 263,  264 

Phosphoric  acid  tests 191 

Phosphorus  pentachloride 192 

Phosphorus 183 

amorphous 186 

and  halogens 191 

and  oxygen 35,  188 

and  hydrogen 186 

bromide 97 

bromides 193 

iodides 193 

pentoxide 83,  189 

red 184 

solubility    1»5 

trichloride 192 

trioxide 189 

Photography 301 

Physical  changes 20 

properties  of  water 59 

Pig  iron 377 

Planets  and  carbon  dioxide 211 

Plaster  cast 317 

of  Paris 317 

Plating  solution 298 

Platinum 390 

alloys 392 

group 389 

stills 160 

Plumbago 203 

Porcelain 341 

Positive  elements 249 

Potassium 45,  281 

bromide 283 

carbonate..  .   282 


Potassium  chlorate 26,  31,  284 

chloride 283 

cyanide 229 

dichromate 71,107,  151 

ferricyanide   386 

ferrocyanide 386 

hydroxide 2*2 

iodide 64,  283 

nitrate 125,  283 

oxidation  by 128 

oxide 281 

permanganate 65,  372 

persulphate 165 

silicate 175 

sulphate 75,  283 

Pottery 341 

Powder  of  Algaroth 358 

Precipitant 14 

Precipitates,  difference  in 16 

Precipitation 11 

by  alcohol 12 

by  reagents 13 

by  water 12 

changes 21 

of  gums  and  resins 12 

of  magnesium  sulphate.  ...  14 

Prefixes 144 

Preparation  of  bromine 94 

of  chlorine 66,  71 

of  hydrochloric  acid 73 

of  hydrogen 46 

of  iodine 99 

of  nitric  acid 125 

of  oxygen 31 

Pressure,  correction  for.. 115 

Priestley 30,  132 

Printing,  silver 301 

Problems  on  gas  volumes 112 

Properties  of  nitrogen 109 

of  water. .              59 


INDEX. 


405 


Proust 195 

Purification  of  water 03 

Purple  of  Cassius 309 

Pyrites   150 

Pyrogallol 308 

Pyroligneous  acid 228 

Pyrolusite 370 

Pyrophosphoric  acid 190 

Pyrosulphuric  acid 162 

Qualitative  relations 29 

Quantitative  relations 29 

Ramsay 117 

Rare  earths 343 

Rate  of  diffusion 49 

of  expansion 112 

Rayleigh 117 

Reaction,  copper  and  nitric  acid  129 

Reactions 22 

analytical 23 

dry 26 

general 24,  25 

in  solutions 26 

of  gases 27 

of  sulphuric  acid 161 

synthetic 23 

Red  lead 350 

phosphorus 184 

Reducing  agents 151 

Reduction  of  copper  oxide 55 

of  nitrates 123 

Reducing  power  of  hydrogen. .  50 

Refining  of  gold 307 

Refraction  of  light 321 

Refrigeration 137 

Reinsch  test 201 

Rhodium 393 

Rose's  metal 360 

Rosin..,  12 


Rubidium 286 

Ruthenium 393 

Rutherford 106 

Safety  lamp 219 

Sal  ammoniac 3 

Salt 66 

Saltpeter 125 

solubility  of 6 

oxidation  by 428 

Salts 104,  257 

Scheele...30,  66,  106,  194,  202,  319 

Schoenbein -. 41 

Schweinfurth  green 198,  294 

Seaweed 98 

Selenium 146,  170 

Sewage,  oxidation  of 131 

Siemens  steel 380 

Silicates 174,  266 

Silicic  acid 173 

Silicon 172 

and  halogens 178 

and  hydrogen 177 

and  oxygen ]  72 

fluorides 103,  178 

Silver 294 

arsenate 199 

arsenite 199 

bromide 300 

chloride 300 

chromate 304,  366 

iodide 300 

nitrate 299 

paper 301 

plating 298 

refining 297 

tests 305 

Silvine 281 

Slaked  lime 316 

Slow  precipitation 15 


406 


INDEX. 


Smelting  process 296,  306 

Soda  fountains 212 

Soda  manufacture 275 

Sodium 44,  272 

amid 138 

bicarbonate 27,  278 

bromide... 279 

carbonate 276 

chloride 72,  279 

hydroxide 274 

hyposulphite 164 

iodide 279 

nitrate , 280 

oxides 274 

phosphates 191 

silicate 175 

sulphate 7,  75,  86,  280 

sulphite 154,  279 

thiosulphate 10,  164,  280 

Soft  water 63 

Soils,  nitrates  in 131 

Solids,  liquids  and  gases 2 

Solubility 5 

of  alum 6 

of  ammonia 136 

of  bromine 96 

of  chlorine 70 

of  phosphorus 185 

of  sulphur 147 

of  sulphurous  oxide 151 

Soluble  glass 175 

Solutions 8 

supersaturated . .    7 

Solution  by  acids 5 

reactions 26 

Solvay  process 278 

Solvent  action  of  water 62 

Specific  gravity 62 

heat 60,  238 

determination  . .            .  2C9 


Specific  gravity  of  compounds,.  267 

weight 62 

Spectroscope 320 

Spectrum 323 

analysis 325 

Spontaneous  combustion 36 

Standard,  water  as  a 61 

Starch 224 

and   iodide,  ozone  test 42 

and  nitric  acid 123 

Stassfurt  mines. 94 

States  of  matter 2 

Steel 380 

Stibnite 356 

Stills 17,  18,     19 

Strontianite 318 

Strontium 318 

Sugar  tests 224 

Sulphantimonates 358 

Sulphantimonites.  . . .- 358 

Sulphates 163,  260 

Sulphide  ores 146 

precipitates 167 

Sulphides 258 

Sulphites 154,  260 

Sulphur 2,  146 

and  carbon 230 

and  chlorine 169 

and  hydrogen 165 

and  niter 128 

and  oxygen 35,   149 

dichloride 169 

dioxide 85 

monochloride 169 

tetrachloride - 170 

trioxide   155 

Sulphuretted  hydrogen 165 

Sulphuric  acids 162 

Sulphuric  acid 156 

plant 158 


INDEX. 


407 


Sulphuric  oxide 153,  155 

Sulphurous  acid 152,  1 54 

oxide H9 

bleaching  by 150 

oxidation  of 153 

uses  of 155 

Supersaturated  solutions 15 

Supersaturation 7 

Symbols  81,  82 

Sympathetic  ink 388 

Synthesis  of  water 55 

Synthetic  reactions 23 

Table,  atomic  weights 80,  81 

composition  of  glass 177 

gas  composition 217 

molecular  weights 242 

of  valency 142 

periodic  system 252 

specific  heats 239 

sulphur,   selenium  and  tel- 
lurium compounds.  ...  171 

Turner's  weights 236 

vapor  tension Gl 

valence  in  groups 254 

weights  of  Berzelius 235 

Tantalum 355 

Tartaric  acid 27 

Tellurium 146,  170 

Temperature  correction 112 

in  combustion 39 

Tension  of  vapor 60 

Terminations 144 

Tests  for  aluminum 312 

for  ammonium 287 

for  antimony 359 

for  arsenic 200 

for  bismuth 361 

for  cadmium 332 

for  calcium .  318 


Tests  for  carbonates   212 

for  chromium 366 

for  cobalt 388 

for  copper 294 

for  gold 309 

for  iron 386 

for  lead 352 

for  magnesium 314 

for  manganese 374 

for  mercury. . .      335 

for  nickel 387 

for  platinum 392 

for  potassium 285 

for  silver 305 

for  sodium 280 

for  sugar 224 

for  thallium 343 

for  tin 248 

for  zinc   331 

Tetrathionic  acid 164 

Thallium 342 

compounds 343 

specific  heat 248 

Theoretical  considerations 139 

Thermometer 60 

Thiocarbonic  acid 230 

Thionic  acids 164 

Thiosulphates 164 

Thomson 78 

Thorium 353 

Tin 345 

chloride 347 

oxides 347 

pyrite 345 

salts...,    347 

sulphides 347 

Titanium 352 

Triads 250 

Trisilicates 267 

Trithionicacid...  .   164 


408 


INDEX. 


Trommer  test 225 

Tungsten 367 

bronzes 367 

Turner's  table 236 

Turpentine 69 

Uniting  proportions 77 

Unit  of  heat 39,     59 

Uranium 368 

Uranium  compounds 368 

Uses  of  prefixes 144 

of  symbols 81 

Valence 140,  245 

and  periodic  system 254 

variation  in 141 

Valency  table 143 

Value  of  gram 62 

Vanadium 355 

Van  Marum 41 

Vaporization 2 

Vapor  tension 60,  61 

Varec 98 

Vermilion 333 

Vinegar 228 

Vital  air 38 

Volume  composition  of  water. .  54 

theory  of  Berzelius 233 

Wash  bottle 34 

Water 50 

as  standard 61 

composition  of 51 

constants. .  .  59 


Water,  electrolysis 52 

gas 217 

glass 175 

of  crystallization 10 

purification 63 

Waters,  natural 62 

Water,  solvent  action  of 02 

Weight,  molecular 82 

White  lead 351 

vitriol 331 

Witherite 319 

Woehler. 337 

Wolfram  bronzes 367 

Wolframite 367 

Wollaston 236 

Wood  ashes 282 

gas 205 

Wood's  metal 360 

Worm  condenser 18 

Wrought  iron 379 

Yeast 226 

Yttrium 336 

Ytterbium 336 

Zero  point 62 

Zinc 308 

arsenide 195 

chloride 330 

hydroxide 330 

oxide 330 

sulphate 85,  330 

sulphite 155 

Zirconium..                    352 


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